Determination of Presence of Adsorption. In cases where a prepeak or postpeak appears, the presence of adsorption will usually be obvious. In other cases the presence of adsorption may not be readily apparent because the adsorption and diffusion processes occur simultaneously and give rise to charge transfer at the same potential. To detect the presexe of adsorption in these cases, it generally requires that the adsorption process contribute significantly to the total current at the scan rates and concentrations studied. Under these conditions there are three quick tests which can be applied. The first, which is strictly qualitative, is to observe the general shape of the current-potential curves. If they show more symmetry than for the uncomplicated Nernstian charge transfer, or increase in symmetry with increased scan rate or decreased concentrations (Figures 2 and 6), there is reason to suspect adsorption. The other two methods involve quantitative measurement and cannot be applied with certainty if other processes (Le., kinetic effects, coupled chemical reactions, etc.) are present. Varying the scan rate, adsorption processes generally cause the experimental equivalent of the current function, i,lCo*& to increase rapidly with increasing scan rate, while i,/cCo* may remain nearly constant. Very few other processes are likely to show similar behavior (16). The last method involves concentration dependence. While i,/Co* is constant for many processes: when adsorption is present an increase in i,; CO* with decreasing concentration usually is observed, possibly leveling off at some constant value for low concentrations, analogous to Figure 5 . However, caution should be employed when using concentration dependence to detect adsorption as one may be in a range where this is not an important factor. Determination of Type of Adsorption. If weak adsorption is known to be present, plots of the type shown in Figures 3, 4, and 5 can be used to determine whether reactant or product is adsorbed, and which isotherm is most likely valid. For
only one species adsorbed, the behavior is easily characterized, while if both are adsorbed, intermediate behavior will be observed and it may be necessary to resort to different electrochemical techniques or additional theoretical calculations, If the adsorption is strong, the position of the adsorption peak can be used to determine if it is the product or reactant which is adsorbed. However, if both species are adsorbed with significantly different values of the free energy of adsorption AGO, it may be necessary to make plots analogous to Figures 17 and 18 to determine the presence of the less strongly adsorbed material. If both species are strongly adsorbed, the behavior will be about the same as if they were both weakly adsorbed, and neither a prepeak or postpeak will be observed, This is because with both strongly adsorbed, the surface is completely covered at all times and the ratio of the surface concentrations, r o / r R , is Nernstian. (The isotherms considered, Equations 8 and 9, do not assume any interdependence of r~and ro,which must exist, and therefore do not adequately describe this case.) Thus, for the adsorption process, the current function becomes independent of mass transfer, while the portion of the current function dependent on diffusion is exactly as for an uncomplicated reversible charge transfer. Addition of these two current functions then gives the total current-potential behavior. These results indicate that a wide range of effects can be observed when adsorption occurs in stationary electrode polarography, and that adsorption must always be considered when using the method in analysis and in studies of electrode kinetics. RECEIVED for review March 24, 1967. Accepted August 18, 1967. Presented in part at the IVth International Congress on Polarography, Prague, July 1966. During the 1965-66 academic year, R. H. Wopschall held a Public Health Service Fellowship. Work also supported by the National Science Foundation under Grant No. G P 3907.
Adsorption Characteristics of the Methylene Blue System Using Stationary Electrode PoIarography Robert H. Wopschall and Irving Shain Chemistry Department, Unicersitj’ of Wisconsin, Madison, Wis. The reduction of methylene blue was studied to test the theory of stationary electrode polarography for the case in which the product of the electrode reaction is strongly adsorbed. A brief investigation of the mechanism of the electrode reaction, using both aqueous ethanol and aprotic solvent systems, indicated that the reduction probably proceeds through successive oneelectron charge transfers, with a very rapid reversible protonation interposed between the charge transfers. The intermediate appears to be more easily reduced @.Eo o 5 1 mV) than the methylene blue. I n spite of this experimental deviation from the model used i n the theoretical calculations, i t was possible to make empirical calculations which permitted correlation of the experimental results with theory, and the isotherm parameters for the adsorbed product (leuco methylene blue) were calculated. Satisfactory agreement with polarographic results was obtained for methylene blue concentrations below 0.2mM, but at higher concentrations, the adsorption of methylene blue itself becomes important, and only qualitative comparisons were made with theory.
To VERIFY the theory of stationary electrode polarography (1) for systems exhibiting the strong adsorption of the product, and to evaluate the applicability of the theoretical correlations, the reduction of methylene blue was studied. The methylene blue system was selected because it had been studied extensively, and direct comparison with Brdicka’s ( 2 ) polarographic results was possible. In stationary electrode polarography, methylene blue exhibits a typical adsorption prepeak on the cathodic scan, characteristic of a system in which the product of the electrode reaction is strongly adsorbed (Figure 1). Using stationary electrode polarography, the system was studied by Mirri and Favero (3), and by Kemula, Kublik, and Axt (4). Mirri and Favero (1) R. H. Wopschall and I. Shain. ANAL.CHEM.. 39, 1514 (1967). (2) R. Brdicka. Collection Czech. Cliem. Commun., 12, 522 (1947). (3) A. M. Mirri and P. Favero, Ric. Sci., 28, 2307 (1958). (4) W. Kemula, Z . Kublik, and A. Axt, Roczniki Chem., 35, 1009 (1961). VOL. 39, NO. 13, NOVEMBER 1967
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studied the adsorption prepeak exclusively, using only cathodic scans. They reported that the prepeak was symmetrical about the peak potential, and that for high concentrations the adsorption peak current was independent of concentration, in accord with the theory for the strong adsorption of product ( I ) , At low concentrations they noted a decrease in the current, and ascribed this to deviations from Brdicka's theoretical model but, in fact, this is expected because the adsorption becomes partially diffusion controlled under these conditions. Kemula and coworkers, using a single scan rate, found that the adsorption peak position is concentration dependent, and that the system is entirely diffusion controlled at high concentrations. In general, their results also were qualitatively in agreement with the explanation given previously by Brdicka. However, the system is further complicated by the adsorption of methylene blue itself. The Langmuir isotherm parameters for the adsorption of methylene blue on mercury have been determined using a spectrophotometric method by Los and Tompkins (5: 6). From their data, using 1M as the standard solution state, the free energy of adsorption AG for methylene blue was calculated to be 6.7 kcal/mole, while from Brdicka's data for the reduction product, leuco methylene blue, AG = 10.9 kcal/mole. Thus, the methylene blue would be much less strongly adsorbed. Some aspects of the previous work also indicated that the electrode reaction might not be the straightforward reversible system required for rigorous comparison with the theory ( 4 , 7, 8). Therefore, it was necessary to carry out a brief investigation of the electrolysis mechanism of metliylene blue prior to studying the adsorption characteristics of the system. EXPERIMENTAL
Instrumentation. The electrochemical experiments in this investigation were conducted using an electronic multipurpose instrument based on the analog computer amplifiers manufactured by G. A. Philbrick Researches, Inc. (Boston), which incorporated the ideas first suggested by Booman (9) and DeFord (10). The actual circuits used were very similar to some of those described by Underltofler and Shain (11). All experiments were performed with a three-electrode cell system. Two circuit configurations were used. For the slower experiments using a pen recorder, the circuit described in Figure 4, Reference l l was used. For faster experiments using an oscilloscope, the circuit shown in Figure 15a, Reference 12 was used. The signal generator which was used to provide a triangular wave form for stationary electrode polarography, and also a single voltage scan for polarography, was similar to Figure 8 in Reference 11. The circuit was modified to furnish a voltage wave form consisting of half a triangular wave, followed by a constant potential equal in magnitude to the peak of the triangular wave by placing a diode between the square wave output and the integrator input (13). This placement of the diode offers a significant improvement over a previous design ~~
(5) (6) (7) (8)
~
J. M. Los and C. K. Tompkins, Can. J . Clzern.,37, 315 (1959). J. M. Los and C. K. Tornpkins. J . Chern. PWys., 24, 630 (1956).
R. C . Kaye and H. I. Stonehill, J . Chem. Soc., 1952, 3244. L. Michaelis, M. P. Schubert, and S. Granick, J. Am. Chent. Soc., 62, 204 (1940). (9) G. L. Booman, ANAL.CHEM., 29, 213 (1957). (10) D. D. DeFord, Division of Analytical Chemistry, 133rd National Meeting, ACS, Sail Francisco, Calif., April 1958. (11) W. L. Underkofler and I. Shain, ANAL.CHEW,35,1778 (1963). (12) W . M. Schwarz and I. Shain, Ibid.,p. 1770. (13) R. H. Wopschall, Ph.D. Thesis. University of Wisconsin, Madison, Wis., 1967.
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ANALYTICAL CHEMISTRY
1.0
0.5 c)
Y
i
5
0.0
U
a: 3
0
-0.5
-I .o
0.0
-0.1 -0.2 -0.3 VOLTS vs. S.C.E.
Figure 1. Cyclic stationary electrode polarograms for methylene blue in buffered solution (pH 6.5; scan rate 44.5 mV,kec) A , 1.00 X 10-*M; B, 0.70 X 10-'M; C,0.40 X 10-'M
(14) where the diode was placed between the output of the multivibrator section and ground. For those experiments using scan rates greater than 100 mV/second, the readout device was either a Tektronix model 536 oscilloscope (Tektronix, Inc., Beaverton, Ore.), with a Type D plug-in unit in the vertical input and a Type T time base plug-in unit in the horizontal input, or a Tektronix model 5 4 5 4 oscilloscope, with a Type D plug-in unit, Both oscilloscopes were equipped with Polaroid cameras. The film was either Polaroid Type 42 or 47 for prints, or Type 46-L for transparencies. Enlargement of the transparencies made it possible to measure areas accurately. For stationary electrode polarography at slower scan rates (less than 500 mV/second, and overlapping the oscilloscopic range), and also for coulometry, a Leeds & Northrup Type G Speedomax pen recorder (Leeds & Northrup Co., Philadelphia, Pa.), which had a full scale response time of 0.4 second and chart speed ranging from 3 to 90 inches/minute was used. For polarographic experiments with the dropping mercury electrode a similar recorder was used, except the full scale response time was 1 second, and the chart speed was 1 inch/minute. Cell and Electrodes. The cell consisted of a 100-ml borosilicate glass weighing bottle provided with a 50/12 standard taper joint, A Teflon lid was machined to fit the top of the cell, with appropriate holes provided to allow insertion of the electrodes, etc. The Pt-wire type HMDE which was used for most work, was constructed by sealing a short piece of platinum wire (26 gauge) in the end of 6-mm diameter soft glass tubing drawn to a point, and polishing the end until the platinum wire formed a smooth flat surface with the glass. The platinum was etched back several thousandths of an inch with aqua ~~
(14) R. S . Nicholson and I. Shain, AKAL.CHEM., 37, 190 (1965).
regia and then plated with mercury. Mercury drops, generally two but sometimes three, from the dropping mercury electrode were collected in the scoop and hung on the electrode to form the HMDE. At high methylene blue concentrations, the solutions were opaque, and hanging the mercury drop became difficult. Under these conditions, a microburet type HMDE electrode (Metroha Ltd., Switzerland) was used. This electrode consists of a mercury reservoir with a piston connected to a capillary, through which a hanging mercury drop electrode is formed. The electrical c a t a c t is made through the mercury contained in the capillary, the piston, and then to the external connection through a bearing and the micrometer threads. The resistance between the piston and the external connection varied from 6 to 10 a,and was neither constant nor reproducible. By inserting a platinum wire through the side of the body of the microburet and sealing it with epoxy resin (one section is Plexiglas, and making this modification is simple), contact was made with the piston through a small amount of mercury which usually remains above the reservoir seal. Measurements showed that the resistance of the modified microburet electrode was constant and equal to the resistance of the capillary, which easily could be reduced to less than 1 a. When mercury drops are removed from the electrode, usually by lightly striking the body of the microburet, there is a tendency for solution to rise a few millimeters in the capillary. Normally this is prevented by applying a water-repellent coating (Desicote, Beckman Instruments, Inc.) so that the capillary is not wetted. However, in some solvents (e.g., DMSO) this technique is not applicable because the water-repellent coating is soluble. If the capillary end is polished carefully, the drop hangs from the mercury column rather than from the glass surface, and there is less tendency for solution to fill the capillary. As received from the manufacturer, surface imperfections near the capillary opening generally caused the mercury drop to hang from the glass, and the shielded area was up to five times that when the tip was polished. Polishing was carried out until imperfections could not be observed under 2 0 X magnification. In some cases where the solutions were not opaque, both types of electrode were used in replicate experiments, and no differences in the polarograms were observed. In either case, the radius of the hanging mercury drop electrode was of the order of 0.06 cm. The reference electrode contained three sections separated by 10-mm fine fritted glass disks. The first compartment was a saturated calomel electrode; the center section contained a concentrated solution of indifferent electrolyte, usually 1M sodium nitrate; and the remaining section, terminating in a Luggin capillary, was filled with the solution to be investigated. The counter electrode consisted of a helically wound platinum wire (approximately 0.5 meter, 26 gauge) sealed through a soft glass tube. The dropping mercury electrode was the conventional type used in polarographic work and had a drop time of about 5 seconds. The deaerator consisted of a glass tube with a coarse fritted glass disk. Normally the deaerator was raised to provide a cover of oxygen-free gas over the solution while carrying out the experiment. When necessary to minimize cell vibration, the entire cell assembly was mounted on a heavy platform supported by an inflated rubber inner tube. Most experiments were carried out at room temperature (25 O i 1 O C). Materials and Chemicals. Prepurified grade nitrogen or commercial grade argon was used for deaeration of the solutions. Traces of oxygen were removed by bubbling the gas through an acidified vanadous sulfate solution. Before entering the cell, the nitrogen or argon was then passed through a 0.1M sodium hydroxide solution to remove the acid, through water to remove the base, and finally through the solvent being used. Deaeration times were 5 minutes or longer.
Laboratory distilled water was further purified by a continuous operation still (15). Ethanol (95 %) was redistilled through a 1-meter column (glass helices) with the middle fraction retained. The dimethylsulfoxide (DMSO) was “Baker Analyzed,” and the dimethylformamide (DMF) was Matheson “Spectroquality Reagent”; both were used without further purification. Methylene blue (Merck U.S.P.) was recrystallized from ethanol-water three times and dried under vacuum over phosphorus pentoxide. From this, a stock solution was prepared, and diluted for all concentrations. Several workers have reported difficulty in determining the number of waters of hydration (values from 1 to 3 have been reported) with the most frequently reported value being two (8). Therefore, in reporting concentrations, two waters of hydration have been assumed, which might possibly introduce as much as a 5 % error. However, when considering the accuracy of the experimental results, this error is negligible. For the methylene blue system in ethanol-water, BrittonRobinson buffer (pH 6.5) was used for the indifferent electrolyte, except for some work at the high scan rates when additional indifferent electrolyte ( 0 . 5 M sodium nitrate) was present. The buffer was prepared using U.S.P. grade barbitol, with all other components being reagent grade. In DMSO, solutions were prepared by dissolving weighed quantities of methylene blue and diluting to volume. Hydrogen ion was provided by the addition of known quantities of 90 HN03 dissolved in DMSO. CHARACTERISTICS OF THE METHYLENE BLUE SYSTEM
Methylene blue is generally available as the chloride salt
CIwhich is readily soluble in both ethanol and in water. In solution it forms dimers, and it may form higher aggregates when the concentration is above 1Ou2M(16). Vetter and Bardeleben (I 7) reported values for the diffusion coefficients which are in good agreement with other results: Dmonomer = 7.6 X cm2jsecond,D d i m e r = 0.95 X cm2/second. They also reported the dissociation constant for the dimer, K = 3.2 X loM4molei’liter. While the dimerization is an important factor in the interpretation of both conventional polarography and stationary electrode polarography, it can be corrected for empirically (18) by using the effective diffusion coefficient at each concentration. There is a possibility that the reduced form of methylene blue also dimerizes, but the evidence is inconclusive. Methylene blue undergoes a two-electron reduction. The reduced form, leuco methylene blue, is a colorless material which is a strong base, and which can be readily re-oxidized to form the highly colored methylene blue-the overall reaction being reversible. (15) K. Hickman, ANAL.CHEM., 36,1404 (1964). (16) P. J. Hillson and R. B. McKay, Trans. Faraday SOC., 61, 374 (1965). (17) K. J. Vetter and J. Bardeleben, 2. Elektrochem., 61, 135 (1957). (18) I. M. Kolthoff and J. J. Lingane, “Polarography,” VoI. 11, Interscience, New York, 1952, p. 845. VOL. 39, NO. 1 3 , NOVEMBER 1967
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The classical study of the redox and acid-base properties of methylene blue was conducted by Clark, Cohen, and Gibbs (19). In general, their results indicated that the amino nitrogen groups are so weakly basic that they are never protonated, while the bridging nitrogen is strongly basic and is always protonated over the p range 1-13. On the other hand, for the leuco methylene ue, the amino nitrogen groups apparently can be protonated along with the bridging nitrogen. Acid dissociation constants for the amino groups have been 0.95 X and K2 = 1.36 reported (at 20" C> as ICl x 10-6 (20). Vetter and ardeleben (17) have proposed overall redox mechanisms which do not take into account these known acid-base properties ; although additional study may be needed, for the present investigation concerning adsorption the acid-base properties were not an important consideration as all experiments were conducted in buffered solution at a single pM. Although Kemula, Kublik, and Axt ( 4 ) were primarily interested in the adsorption phenomena, they did conduct some studies under conditions where the system was exclusively diffusion controlled (high methylene blue concentrations and acid media). Under these conditions they found the anodic and cathodic diffusion peak potentials differed by 35 mV, and considered this to be in agreement with the theoretical value for a two-electron reversible reaction. In fact, one would not expect a 35-mV difference, but about a 28.5-mV peak separation for a two-electron reversible reduction, or about a 57-mV separation for a one-electron reduction (21). Thus, their results indicate that the reduction mechanism for methylene blue may not be an uncomplicated two-electron, reversible charge transfer, but rather might invole a one-electron reduction intermediate as proposed by Michaelis et al(8). INVESTIGATION OF THE METHYLENE BLUE REDUCTION MECHAWESM
There are several mechanisms which can account for the anodic-cathodic peak potential separation of 35 mV observed by Kemula et al. (4). Three of the possible mechanisms would be: charge transfer kinetics (quasi-reversible charge transfer), a chemical reaction coupled between two charge transfers (ECE mechanism), or two successive oneelectron charge transfers (EE mechanism). The ECE and EE mechanisms both require an intermediate reduction state between methylene blue and leuco methylene blue. To test for each of these possible mechanisms, studies of methylene blue were conducted in both buffered ethanol-water (pH = 6.5) and in nonaqueous, aprotic solvents. Studies in E t ~ a ~ ~ ~ - WFor a ~ ethe r ~mechanism studies, stationary electrode polarograms were obtained for methylene blue at concentrations where adsorption effects are relatively unimportant (3mM). The peak current measurements were made over a range of scan rates from 15 mV/second to 20 Vlsecond. If the charge transfer were quasi-reversible, the current function would decrease and the peak separation would increase with increased scan rate (21). But the data indicated that the function (i,)c/.\/; [where (i,), is the cathodic peak current and I; is the scan rate] was nearly constant to
(19) W. M. Clark, €3. Cohen, and H. D. Gibbs, Public Health Rept. (US.), 40, 1131 (1925). (20) B. P. Nikol'skii. M. S . Zakhar'evskii, and V. V. Pal'chevskii, Uch. Z a p . Lenitigr. Gos. Utiilie., Ser. Khim. Nank. No. 15, 26 (1957); [Clzem. Absrr., 5 2 : 9812 (1958)l. (21) Ea. S. Nicholson and I. Shain, ANAL.GNEM.,36, 706 (1964).
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10 V/second and then increased rather than decreased as would be expected- for the quasi-reversible case. (This increase in ( i & d u at high scan rates was due to adsorbed methylene blue.) In addition, the ratio of anodic to cathodic peak currents, (ip)a/(ip)G, was close to unity for scan rates less than 10 Vlsecond, indicating that the methylene blue system behaves reversibly. A decrease observed at higher scan rates again could be attributed to the adsorption of methylene blue. Careful measurement of the difference between the peak potential and the half peak potential ( E p - E P / Jgave a value of 34 mV at the lower scan rates, in good agreement with the results of Kemula et al. ( 4 ) . With increased scan rate, only minor broadening of the peaks and slightly increased peak separations were observed, and this was attributed to uncompensated iR drop. If a slow charge transfer were involved, the peaks would broaden and their separation would increase markedly with increased scan rates. On the basis of these results, it was concluded that the observed peak separation could not be accounted for by charge transfer kinetics. The results also seem to indicate the mechanism cannot be an ECE mechanism in which the chemical reaction is irreversible, because the theory presented previously (22) indicates that for this type of system the function (iP)J& must decrease with increasing scan rate, while the methylene blue shows either a constant or increasing value for this parameter. However, the reaction could be an ECE mechanism with a fast, reversible chemical step. For the third possible mechanism, the EE mechanism, previous calculations (23) had indicated that with identical reduction potentials for the two steps, a peak separation of about 42 mV would result, but with the shape essentially unchanged from the Nernstian behavior for a single charge transfer. Also, for an EE mechanism, the current functions [or ( Q c / d u T r (ip)a/(ip)c] exhibit no scan rate dependence. Thus, an EE mechanism could account for peak separations intermediate between that expected for a one-electron and that expected for a two-electron charge transfer. If the methylene blue undergoes this kind of a two-step reaction, the reduction potentia1 for the intermediate must be more anodic than the reduction potential for methylene blue because the methylene blue system appears to be even more like a two-electron charge transfer than two successive charge transfers with identical reduction potentials. Using results of Polcyn (24) theoretical calculations for the EE mechanism gave stationary electrode polarograms which agree with the methylene blue system if the reduction potential, E", for the intermediate is 51 mV anodic of the reduction potential for methylene blue. Because of the ease of reduction of the intermediate, it would never be present in quantities greater than about 17z of the total material, in agreement with the proposal of Michaelis et al. (8). Studies in Aprotic Solvents. The characteristics of an EE mechanism and those of an ECE mechanism with a fast reversible chemical reaction are identical, provided the experimental times are less than the half-life of the chemical reaction. Thus, it appeared possible that the electrode reaction could involve a fast protonation coupled between the two charge transfers, To investigate the possible existence of an intermediate, experiments were carried out in two aprotic solvents, D M F and DMSO. In preliminary studies similar (22) Ibid., 37, 178 (1965). (23) D. S. Polcvn and I. Shain. Ibid., 38,370 (1966). (24j D. S . Poicyn, University of Wisconsin, Madison, Wis., unpublished data, 1965.
6
ir
42-
u
h
0-
-2
c--
-
-4c
POTENTIAL Figure 2. Polarograms of methylene blue in dimethylsulfoxide 1.OmM methylene blue chloride (except curve 0).A , no acid; B, approximately 0.5mM nitric acid; C, approximately l.OmM nitric acid; D,methylene blue nitrate; E, hydrogen wave with excess acid
behavior was observed in both solvents, and for convenience, DMSO was used as the solvent in most of the nonaqueous work. Conventional polarograms were obtained on methylene blue in DMSO with potassium nitrate as the supporting electrolyte. The chloride salt of methylene blue gave three waves as shown in Figure 2 , curve A . The anodic wave was identified as due to chloride and was not as well defined as either of the two cathodic waves. The anodic wave could be eliminated if the nitrate salt of methylene blue was used in place of the chloride salt (Figure 2 , curve 0). The first (I) and second (11) cathodic waves correspond to the one- and two-electron reduction of methylene blue, respectively. Coulometric reduction in the region of wave I gave a value for n slightly greater than one, while in the region of wave I1 a value of n of 2.02 was obtained. The two half wave potentials are separated by 738 mV, indicating that a reasonably stable intermediate is formed by the one-electron reduction of methylene blue. The intermediate, appearing deep red in color, was easily air-oxidized back to the original blue. The final product (two-electron reduction) was orange-yellow and was also easily air-oxidized to the original blue. Addition of excess acid to the intermediate caused it to turn green immediately, and again air oxidation gave the original blue. Waves I and I1 were of nearly the same height in the absence of any additional hydrogen ion (there is a small amount due to water in the DMSO). Addition of water showed some effect in distorting the wave shapes, but the wave heights remained nearly constant. However, addition of hydrogen ion caused wave I to increase while the total current in the region of wave I1 remained nearly constant (curves B and C). The increase of wave I was nearly linear with the addition of acid until it blended almost completely with wave I1 to give a single wave. As the acid concentration was further increased, a hydrogen wave appeared cathodic of wave I1 (curve E ) and a nonreproducible maximum (which has not been shown) appeared on the rising portion of wave I. Using a microburet type electrode, stationary electrode polarograms were obtained on the nitrate salt of methylene blue. The scan rate was varied from 100 mV,kecond to 100 V/second with no appreciable change in peak shape and position except for an apparent iR drop effect. Two cathodic and two corresponding anodic waves were observed, all essentially the same height, as shown in Figure 3. Both reactions ap-
I
POTENTIAL Figure 3. Stationary electrode polarogram of methylene blue in dimethylsulfoxide
CO*= 3mk1, L'
=
330 mV/sec
peared to be one-electron reversible charge transfers with about 750 mV separation, or essentially the same as the polarographic half wave separation. These results indicate that in DMSO, methylene blue (deep blue) is reduced in a one-electron charge transfer to form a deep red intermediate. This intermediate is unstable when protonated, and either immediately reacts at the electrode surface (another one-electron charge transfer) to form the twoelectron product, leuco methylene blue (yellow-orange), or disproportionates to form methylene blue and leuco methylene
POTENTIAL Figure 4. Stationary electrode polarogram of methylene blue under conditions where adsorption effects are relatively minimized Concentration, 0.15mM; scan rate, 10.5 mV/sec. Points : experimental; line : theory for uncomplicated reversible case with n = 1.66, and also theory for two successive one-electron transfers with the intermediate more easily reduced than the reactant, AE' = 51 mV. These two theories show differences which are too small to be noted on this plot VOL. 39,
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the theoretical calculations, and it was assumed that this procedure could account for the presence of the intermediate in the EE mechanism. There is no doubt about this being a valid procedure in describing the diffusion-controlled peak, but unfortunately, in the presence of adsorption, there is no assurance that the procedure is also valid in describing the adsorption prepeak. However, the prepeak showed no major deviations from reversibility, no split waves were observed, and the anodic and cathodic currents were essentially identical. Thus it was not unreasonable to assume that the adsorption process would affect the two individual charge transfers to the same extent, and that the difference in Eo’s for the successive steps would not be markedly changed. On this basis, it was felt that no major errors in calculating the adsorption isotherm parameters would be introduced. ADSORPTION IN THE METHYLENE BLUE SYSTEM
-0.1 -0.2 VOLTS vs. S.C.E.
-0.3
Figure 5. Stationary electrode polarograms of methglene blue at two scan rates Concentration = 0.1 mM, pH 6.5, 50 wt A , 82.4 mV/see; B, 22.4 mV/sec
ethanol-water.
blue. A disproportionation probably takes place on addition of acid to the deep-red solution, giving a mixture of methylene blue and leuco methylene blue, which possibly accounts for the green color. The Reduction Mechanism-Summary. Consideration of these results indicated that the reduction of methyIene blue in ethanol-water is probably the result of two successive one-electron reversible charge transfers, with a fast protonation of the intermediate to form a species which is more easily reduced than methylene blue. To obtain results for a theoretical model which combines both the EE mechanism (or the equivalent ECE mechanism), with the strong adsorption of product would be difficult, and presentation of results would be complex. Therefore, to use methylene blue for correlation with theory, some simplifying assumptions had to be made. For an EE mechanism, in which only a single wave is observed, the shape of a theoretical stationary electrode polarogram is the same as the shape of a single reversible charge transfer. Thus, it is possible to fit theory for a reversible charge transfer to an experimental EE mechanism simply by using a nonintegral value for n to normalize both the peak height and the potential spread of the wave. For methylene blue, a value of n equal to 1.66, rather than the coulometric value of 2.0, gave a wave of the appropriate shape. Figure 4 shows that the two theoretical plots, for a reversible charge transfer with n = 1.66, and for an EE mechanism with a 51mV reduction potential separation are essentially the same, and that the theoretical curves agree reasonably well with the experimental curves for methylene blue, except for the foot of the wave where some error exists because of the adsorption of product. Thus, to make correlations between experiment and theory for the methylene blue system, a value of n = 1.66 was used in a
ANALYTICAL CHEMISTRY
Stationary electrode polarograms of methylene blue in 55 wt ethanol-water buffered at a pH of 6.5 with BrittonRobinson buft’er, and in some cases containing an additional 0.5M sodium nitrate, were obtained at scan rates ranging from 5 mV,kecond to 16 V/second and over a concentration range of O.OlSmA4 to 1.5mM. For each concentration, the appropriate scan rate range was selected so that both peaks were measurable, normally restricting the ratio of adsorption peak to diffusion peak to the range between 0.1 and 1.5. As indicated by the theoretical study of strongly adsorbed product with weakly adsorbed reactant ( I ) , two concentration regions can be considered: low concentrations where the adsorption of reactant has small effect on the current-voltage curves, and high concentrations where the adsorption of reactant must be considered. For concentrations below about 0.2mM, the methylene blue system fits the simpler theory in which the adsorption of reactant is ignored. This concentration range was considered in detail and the adsorption isotherm parameters were determined. However, above 0.2mM the adsorption of methylene blue must be considered. This concentration range was considered qualitatively to show that the behavior was similar to that predicted by theory. Low Concentrations--Co* Less than about 0.2mM. A qualitative investigation of the methylene blue system was carried out to confirm that the system followed known trends with variation of the methylene blue concentration and scan rate. The cyclic polarograms (Figure 1) show a cathodic prepeak and an anodic postpeak which are nearly equal in magnitude and shape (roughly symmetrical about the peak potential), as required by theory. However, the peaks are slightly shifted, with the anodic postpeak appearing about 10 or 20 mV anodic of the cathodic prepeak. This separation couid indicate that a rate-controlled adsorption-desorption process may be involved. But the shift is small and considering the other approximations which were made, this deviation couid be ignored, The polarograms in Figure 1 also show the trends predicted from theory, that with an increase in concentration, the adsorption peak current (not current function) should remain nearly constant, while the diffusion peak current should increase. At any particular concentration where both peaks are observed, for an increase in the scan rate, the theory predicts an increase in the adsorption peak current function, with a slight decrease in the diffusion peak current function due to depletion of material near the electrode. This behavior for the methylene blue system is shown in Figure 5 and again is in qualitative agreement with the theory. To determine the adsorption isotherm parameters, a quanti-
I O
0
/
10-4
10-2
10-3
cc,/V'l2 Figure 7. Ratio of (ip),/Co*v'G to its value as Co*/u 4 a,as a function of C o * / d % (moles-1-1-u-1~2sec1'2) Points: experimental; line: theory for UEF/RT= 1.6 X 1Q-2mV-1
0.0 0
2 Jr/C,:
4 I
I
/O
A
I /
Figure 6. Ratio of adsorption peak to diffusion peak, as a function of z/U/CO* (cl'2-sec-1'2-liter-mole-1 X 10-3)
,,
0.8 I
7
CO*ranges from 0.015 to 0.15mM,u ranges from 5 to 600 mV/sec
tative correlation was made between experiment and theory. This involved plotting p against d;/Co*, and (ip)c/Co* against Co*/du-for a variety of scan rates and concentrations. [The notation used is the same as defined previously ( I ) ] . The first of these correlations is shown in Figure 6 where the plot is linear, in agreement with theory, and the scatter is not more than would be expected. These results, which include concentrations as high as 0.15mM, do not exhibit the nonlinear behavior shown in Figure 20 of Reference l . Thus, ignoring the adsorption of methylene blue in the correlations is valid for concentrations below 0.1 5mM. The second plot of ( i , ) c / C ~d *Fm. Co*/.\/U? (Figure 7) is actually normalized to plot as the ordinate the ratio of the parameter (i,>,/cO* d i t 0 its limiting value as ~ o * / --+ v ,in accord with the theoretical plot. There is a large amount of scatter in the data, probably because of a combination of experimental error and deviations from the theoretical model. However, the general trend is in sufficient agreement with the theory to allow one to use the theoretical model to calculate the isotherm parameters. Using the correlation shown in Figure 10 of Reference I , the value of aRF/RT was calculated from AE, for each stationary electrode polarogram. Then, the value of p was used to obtain an estimate of PR from Table I1 of Reference I from which a second approximation for URFIRTwas determined using the data presented in Figure 9 of Reference I . The average value for this parameter was (1.6 i 0.4) X lo-? mV-1 for leuco methylene blue. The value for PPR was 0.360. Comparison of the theoretical value of ~ ( F Rwith the experimental slope of the p us. ~ ? C O plot * (Figure 6) makes it possible to calculate the parameter 4rRs l/nF/RT.lrDo. Assuming n equals 1.66, the value of r R B / which was obtained was 0.368 X ( m o l e - s e ~ o n d ~ ~ ~ ) / Assuming cm~. an average
6
dz
Q
0.4
I /
0.0 0
I
I
I
1
2
3
v"2/c; Figure 8. Ratio of adsorption peak to diffusion peak for methylene blue as a function of GjC,,* (ul;2-sec-l/2-mrPle-'-liter) Experimental points: 0,0.4 mM; 9, 0.7mM;
c),
1.OmM
Theoretical lines A , B, and C correspond to 0.4,0.7, and 1.0 miM methylene blue, selected to give best fit at 0.4mA4, with P E = (1.42 X 10-s)P~and V R = 1.26 PO. Dashed line is low concentration behavior
value for DOof 6.2 x 10-6 cm2/second, one obtains a value of (9.1 j=0.8) x 10-11 molejcm? for rRB, or about half the value which was obtained by Brdicka (1.62 X 10-10 mole,km2). However, measuring the area under the adsorption prepeak gives a value of (2.18 + 0.08) X 10-lO moles/cm2 which is about 30% higher than the value reported by Brdicka. The value for rRsfor leuco methylene blue has also been deterVOL. 39, NO. 13, NOVEMBER 1967
* 1533
mined by Lorenz and Schmalz (25) using chronopotentiometry, and found to be 1.5 to 2.0 X 10-lo mole/cm2. It would be expected that stationary electrode polarography would give a high value if charging current were not properly subtracted, while chronopotentiometry might give a low value because of desorption. Considering the accuracy of the three methods of measurement, all results are in agreement and stationary electrode polarography techniques can be used to determine r R s . Once the value of 4 r’; dnF/RTnDo is known, the value of pB can be determined from Co* and u for each polarogram. Then, the value of P R ‘ for each scan rate and concentration was calculated for each polarogram, using the value of AE, (difference between diffusion peak potential and adsorption peak potential) with Equation 45 and Table III of Reference 1. Assuming r R S / l / D R = 3.68 x 10-8, a value of log KE‘ equal to -7.6 i 0.3 was obtained. From this and the prepeak potential, the value of AG” was found to be 9.2 kcal/mole, compared with a free energy of adsorption of 10.9 kcal/mole obtained by Brdicka. High Concentrations--Go* Greater than about 0.2mM. Using high methylene blue concentrations where the adsorption peak is negligible compared to the diffusion peak, it was possible to observe the influence of the weak adsorption of methylene blue on the stationary electrode polarograms. This was particularly evident at high scan rates where marked deviations from reversible behavior were observed. These deviations correlate well with the theoretical calculations for a system involving weak adsorption of reactant as shown in Figures 3 and 4 of Reference I , and are in agreement with the results of Los and Tompkins (5, 6). A semiquantitative correlation of experimental results with theory ( I ) was carried out for data obtained in the concentration range where adsorption of both reactant and product is important. A comparison of the experimental and theoretical results is shown in Figure 8, where the theoretical parameters were selected to fit the methylene blue-leuco methylene blue system at the lowest concentration included. As the concentration is increased, the deviations from linear behavior increase in the methylene blue system just as in the theoretical case. However, these deviations are larger than can be accounted for by theory, probably because the interdependence of the adsorption isotherms for reactant and product was not considered. The ratio of the observed diffusion peak to the diffusion peak where adsorption is negligible is shown in Figure 9 as a (25) W. Lorenz and E. 0. Schmalz, 2.Elektrochern., 62,301 (1958).
1534
0
ANALYTICAL CHEMISTRY
\
1.4c
1.2 &---
I
I
0 log cc;
/V”V
Figure 9. Ratio of (iJC/l/G to its value as Co*/o a function of Co* Y E (mole-liter -1-sec1’2-c-1’2)
-+ 03 as 9
Experimental points: 0, 0.4mM; 0 , 035mM; 6, 0.lOmM methylene blne Theoretical lines A , B, and C correspond to 0.4, 0.25, and 0.lOmM methylene blue, selected to give best fit at Q.lOmM, with PR = (1.42 X 10-3)Po, and P R = 1 . 2 6 ~ ~ .
-
function of cO*/ d c for various concentrations. Comparison of the theoretical curves (where parameters were selected to fit the methylene blue system at the lowest concentration shown) with the experimental points also shows the similarity of the trends of the theoretical and experimental plots, but quantitative agreement is again poor. While it would be interesting to see if the methylene blue would exhibit a maximum in its current function ratio like Curves A and B, of Figure 17 in Reference 1, it was not possible to obtain sufficiently high scan rates experimentally. While this case involving adsorption of both product and reactant has been studied only briefly, the agreement between theoretical trends and experimental results indicates that the model chosen is a fair approximation. If rate controlled processes and interaction of the adsorption isotherms could have been considered, it is probable that much closer agreement would have been observed between the experimental results and the theoretical calculations.
RECEIVED for review March 24, 1967. Accepted August 18, 1967. Presented in part at the IVth International Congress on Polarography, Prague, July 1966. During the 1965-66 academic year, R. H. Wopschall held a Public Health Service Fellowship. Work also supported by the National Science Foundation under Grant No. G P 3907.
