Thin-Layer Electrochemical Technique for Monitoring Electrogenerated Reactive Intermediates Richard L. McCreery Department of Chemistry, Ohio State University, Columbus, Ohio 432 10
When a thln-layer optlcally transparent electrochemlcalcell is used to generate reactive materlais, the course of the reactlon may be monltored by conventional spectrophotometry. Slnce the solution wlthln the cell Is homogeneous on a ca. 2 0 4 time scale, classlcal klnetic analyses may be applied to the absorbance vs. tlme results, as opposed to the more complex translent electrochemlcal approaches. Furthermore, the Nernstian equllibrium between the electrode and homogeneous solution may be utlilred to greatly increase the apparent half-life by adjustlng the oxldired/reduced (oxlred) ratio. The method has been applled to p-qulnone lmlne hydrolysis, and Its promlse for appllcatlons to more complex reactlons Is dlscussed.
Modern electrochemical techniques have been used extensively to monitor the kinetics of homogeneous reactions following charge transfer a t an electrode surface ( 1 ) . Chronopotentiometry (2), chronoamperometry (3), and cyclic voltammetry ( 4 ) have been particularly valuable for determining the mechanisms and rate constants for reactions of electrogenerated materials. In most cases, the observed current or potential is related to the desired kinetic parameter by a working curve derived from Laplace transform analysis or computer simulation techniques (5). Although these methods have proved very useful for relatively simple kinetic schemes, problems arise with more complicated mechanisms, such as those involving disproportionation or other secondorder reactions of the electrogenerated intermediate. An intensively studied example of such problems is the addition of nucleophiles to aromatic cation radicals (6).The interpretation of kinetic data for this system was frustrated by the fact that working curves for electrochemical response were too similar to allow unequivocal assignment of the mechanism. Much of the reason for this ambiguity is the nonhomogeneity of the solution near the electrode caused by diffusion. The complex concentration profiles which result greatly complicate the kinetics of concentration dependent kinetic schemes. The additiqn of an optical probe to the conventional electrochemical experiment greatly clarifies many kinetic schemes, since one may selectively monitor certain reactants or products (7). Even with spectroelectrochemical monitoring of electrogenerated materials, however, one is dealing with a nonhomogeneous solution, and problems with concentration profiles still arise. Except in certain special cases (8), the spectroelectrochemical response does not allow assignment of complex reaction mechanisms. These problems would be greatly diminished if reactive materials could be electrogenerated in an environment where the solution remains homogeneous on the time scale of a kinetic run. Consider an experiment carried out in a thin-layer electrochemical cell, using an optically transparent grid electrode of the type described by Heineman and Murray (9). Rapid diffusional mixing within the cell assures not only a homogeneous solution, but also provides for establishment of a Nernstian equilibrium between the solution and the 206
electrode (10). After the electrode is used to generate reactive materials, the ensuing reactions can be monitored optically. Two principal advantages result from such an experimental arrangement, provided the mixing time in the cell is much shorter than the duration of the kinetic run. First, the concentration vs. time behavior of solution constituents can be analyzed by simple classical kinetic methods, allowing the reaction order to be determined directly, and avoiding the more complex and ambiguous working curves. Second, the applied potential can be used as a variable which will affect kinetic results. It will be apparent from the discussions below that the potential dependence of rate constants is a valuable diagnostic tool for defining mechanisms. The work described herein was undertaken to establish the utility of this approach for monitoring the kinetics of reactions of electrogenerated reactive materials. To establish the viability of this thin-layer kinetic method the thoroughly studied hydrolysis reaction of p -quinone imine was used as a test system (11,12). This system involves the irreversible pseudo first-order hydrolysis of quinone imine (QI) electrogenerated from p-aminophenol (PAP) to form p-quinone (Q). Since the solution is in the equilibrium with the electrode, the Nernstian ratio of QI to PAP will be fixed by the potential. If the QI/PAP ratio is small, only a fraction of the total PAP is in the reactive form, so the observed reac-
OH
0 NH
0
0
0
t,ion will be slowed down as the potential decreases. This potential dependence can be used to change the time frame of the kinetic run, a quality of great utility for fast reactions. The potential dependence of the observed rate is derived below.
THEORY For a reaction of Equations 1 and 2 in a thin-layer cell, PAP
&I
+ 2e-
(1)
the Nernstian ratio R will be a constant, provided reaction 2 is slow relative to the mixing time in the cell. R will be determined by the difference between the applied potential (Eapp)
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
T a b l e I. Kinetic Data for Quinone Imine Hydrolysis in
1.81 M HzS04 Eapp
0.530 0.540 0.550 0.555 0.560 0.570 0.580 0.580 0.590 0.590 0.60 0.65 0.70 0.75 0.80
R
kobsd
0.142 0.311 0.677 1.000 1.476 3.218 7.015 7.015 15.29 15.29 33.33 1640
x io3
ki(caicd)
2.62 2.54 2.55 2.40 2.65 2.56 2.40 2.78 2.74 2.40 2.53 2.54 2.48 2.59 2.62
0.326 0.602 1.03 1.20 1.58 1.95 2.10 2.43 2.51 2.25 2.46 2.54 2.48 2.59 2.62
m m
m
x io3 -20
-3.0 0
Curve A, EaPp = 0.75 V vs. SCE; Curve B, EaPp = +OS55 V
T h e usual rate law applies t o t h e generation of quinone:
-d'Q1 - k 1[QI] dt
To derive the time dependence of &I, t h e species being monitored here, it is necessary only t o consider a mass balance expression o n total aromatic species:
Making use of t h e definition of R ,
(&)
kl[QI]
(4)
Regardless of t h e magnitude of t h e Nernstian ratio R , one would expect a first-order decay for &I. T h e observed rate constant is related t o the ordinary pseudo first-order constant b y Equation 5 , where R is given b y Equation 3. D
T h e same observed rate constant is obtained if PAP is monitored, although the absolute rates differ by t h e factor R. With t h e PAP system, it is difficult t o find a wavelength where only &Iabsorbs. A t 270 nm, both &I and PAP absorb, although t h e molar absorptivity of QI is some 10 times that of PAP. Despite this complication, t h e slope of the In (Absorbance) vs. time plots will equal kobsd, as shown below.
A = @Op
b[PAP]
In A = In -b
+
+ E@"
6 8 1 0 1 2 1 4 T i m e , minutes
Figure 1. Absorbance at 270 nm vs. time plots for quinone imine hy-
and t h e formal potential (E"') for PAP oxidation; according t o Equation 3
-
4
drolysis in 1.81 M H2S04
2.56 f 0.11 Re1 std dev = 43%
-d[Q1l dt
2
b[QI]
b ) 9 In [&I]
Provided R is a constant, t h e In A vs. time plot will have t h e same slope as a In [QI] vs. time plot.
EXPERIMENTAL Cells. Thin-layer optically transparent cells were constructed from Ylc-in. quartz plate using the design of Heineman (10). A single layer of 3-mil Teflon spacer yielded a cell thickness of ca. 0.015 cm. The choice of cell thickness in any thin-layer work involves a trade-off of optical path length and IR drop vs. the diffusional mixing time. In this case, the molar absorptivity of &I is large enough and the rate con-
stants slow enough that neither path length nor mixing time caused observable error, To suit a particular application, the thickness could easily be altered, provided the trade-off is kept in mind. Gold mesh, 200 wirelin., served as the working electrode, having an area within the cell of 2.0 cm X 1.8 cm. The relatively large electrode area was used to minimize edge diffusion during the longer kinetic runs. Vinyl tape was used to mask off all but the center 1mm X 3 mm area of the electrode, and the gold mesh was placed within 2 mm of the bottom of the cell. The reference and auxiliary electrodes were placed in a 5-ml cup below the thin-layer cell. More sophisticated arrangements to avoid ohmic potential losses were unnecessary. The entire cell arrangement was placed in the argon filled sample chamber of the spectrophotometer. Spectrophotometer. Absorbance measurements were made with a single beam, fixed wavelength spectrophotometer based on a Beckman DU prism monochromator and hydrogen source. The output of the photomultiplier (EM1 6256) was amplified by an operational amplifier circuit and recorded on a log nltime strip chart recorder. Despite single beam operation, drift was not significant enough to cause observable error during a kinetic run. The beam emerging from the DU monochromator is sufficiently narrow that the cell could be placed directly between the exit slit and photomultiplier, with no other optics required. Potentiostat. A kepco 50-watt power amplifier served as the control amplifier of a conventional DeFord-type potentiostat. The only special requirement of the design was accurate (flmV) potential control, but in practice the potential was set with the aid of a high accuracy digital voltmeter. All potentials are expressed relative to the saturated calomel electrode (SCE). Reagents. p - Aminophenol was obtained from Matheson, Coleman and Bell, and was used without purification. Solutions were prepared with standardized sulfuric acid and degassed thoroughly before use. Procedure for Kinetic Runs. After the cell was filled and positioned in the sample chamber, it remained only to apply the potential and monitor the absorbance at 270 nm. The success of the approach described here depends on the rapid establishment of a constant ox/ red ratio in the optical path. In addition to diffusional mixing time, two phenomena slow down the establishment of R ; the charge transfer rate may be slow, and IR drop may result in a nonhomogeneous potential. It was found empirically that "overdriving" the system by temporary application of a potential larger than the control value resulted in much faster establishment of the oxlred ratio. The potential was returned to the control value after the approximate oxlred value was achieved, judged by the absorbance. A slightly high potential will help overcome both problems by increasing the charge transfer rate and by partially overcoming IR drop. It is true that this practice will result in some nonhomogeneity of the oxlred ratio, but this effect caused no observable deviation from theory, as shown later in Table I.
RESULTS T h e natural log of absorbance vs. time plot for a n experim e n t carried o u t at +0.75 V in 1.81 M HzS04 is shown i n Figure 1,curve A. T h i s potential is m u c h larger t h a n EO' for t h e PAP/QI couple, so complete conversion t o &I occurs, and one is monitoring the usual first-order decay of &I. Notice that 2 min are required t o reach t h e first-order decay line; this delay is attributed t o t h e time necessary t o carry out the initial electrolysis. T h i s period may be reduced t o less than 30 s with ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
207
-
00*51 7
0020
Figure 2. Observed rate constant vs. applied potential in 1.81 M
Figure 3. Plots of Equation 6
H2S04 Solid line is calculated from Equation 5 using E" = +0.555 V, kl = 2.56X lop3 s-', Circles are experimental points
proper cell design, but i t is important to note that the decay of reactive species must occur in a period much longer than the time required for electrolysis. The uncertainty in the starting time of the run is not important for first-order reactions, but will be of consequence for higher order mechanisms. The value of the slope observed a t f0.75 V is the ordinary first-order rate constant. If the control potential is reduced to a value equal to the formal potential, the plot of Figure 1, curve B, results. Note that the system still obeys first-order kinetics, but that both the initial absorbance and the slope are lower. As predicted, kobsd a t E"' is half the value at +0.75 V. The experiment was repeated with control potentials between 0.520 and 0.8 V, and the graph of Figure 2 was constructed. The open circles are the slopes of the In (Absorbance) vs. time plots; the solid line was calculated from Equations 3 and 5 for an Eo' = 0.555 V and h1 = 2.56 X 10-3. It should be noted that the observed half-life for the reaction carried out at +0.52 V was some 13 times that a t +0.7 V. Rearrangement of Equations 3 and 5 leads to the following expression, which corresponds to the reversibility plot for a poloragraphic wave
A plot of EappVS. log ( k l - hobsd)/hobsd should have a slope of -0.0591/n volts and an intercept equal to E"'. Such a plot for the experiments in 1.81 M HzS04is shown in Figure 3. The slope is 0.0291 V, corresponding to the expected two-electron value. The points a t the extremes of the line were not used in the least squares analysis, and the points a t +0.58 and +0.59 V are averages of the two rate constants determined a t those potentials. An indication of the precision of the method may be obtained by examining the data of Table I. The values for kl were calculated from Equation 5, using the observed rate constants and an E"' value of +0.555 V. The relative standard deviation of the results is 4.5%. The relative standard deviation for other experiments was more typically in the range of 6-8%. In 1.01 M H2S04,a mean h l value of (6.98 f 0.52) X was observed, corresponding to a relative standard devition of 7.4%.
DISCUSSION When the applied potential is more than 0.1 V positive of EO', the ox/red ratio is very large, and the solution is virtually completely electrolyzed to quinone imine. A major advantage of the method in this potential region is simplicity and clarity of data analysis. Classical kinetic plots may be used to determine reaction order, information which is not unambiguously obtained from many other electroanalytical techniques. In addition, electrogeneration of reactive materials provides a 208
for data of
Figure 2
Least squares slope = 0.0291 V, y-intercept = 4-0.555V
convenient means to initiate the reaction. An obvious disadvantage of the method, when used a t potentials much greater than E"', is the relatively slow time frame of the kinetic run. The time required for complete electrolysis limits the method to reactions with half-lives greater than ca. 30 s when the applied potential is on the plateau of the kobsd vs. E plot. When the reaction is conducted at potentials in the region of E"', two additional advantages are apparent. First, the observed half-life can be increased greatly by lowering the potential, since a smaller fraction of PAP is in the reactive form. Near the foot of the h&sd vs. Eappcurve, it should be possible to increase the half-life by an order of magnitude for every 30-mV reduction in Eapp.Despite the 5-10 minute time frame of the technique, i t should be possible to monitor reactions with real half-lives of much less than a second. The limitations of this decelerating effect are presently being determined. Although an increase in half-life of several orders of magnitude should be attainable, i t is unlikely that this technique will be suitable for the submillisecond reactions observable by transient electrochemical methods. The second advantage of the potential-dependent mode is that potential control may be used to selectively generate particular reactive species. For example, nucleophilic addition reactions to oxidized aromatic hydrocarbons are complicated by the fact that two reactive species may be present, the radical cation or dication of the aromatic system. If the solution potential could be controlled, as with this thin-layer method, the distribution of cation and dication would be known, and the kinetics of the reactions of these materials with nucleophiles would be greatly simplified. Marcoux (13)has pointed out the diagnostic utility of potential-dependent chronoamperometry experiments for examining complex mechanisms. The same diagnostic conclusions should be available from the potential-dependent rate constants obtained here. Most electrochemical techniques involving diffusion cannot accurately measure rate constants of reactions with half-lives greater than about 30 s, because of convection in the solution. Thus most previous work on $1 hydrolysis was performed under conditions where k l is much faster than that determined here, and comparison of results is difficult. Using thin-layer reverse current chronopotentiometry, Christensen s-l in 1.02 and Anson ( 1 4 ) found k.1 to be (7.4 f 0.2) X M HzS04. This value is in reasonable agreement with the value of (6.89 i 0.52) X 10-3 found here. The reasons for the slightly poorer precision of the present work are not presently known. When examining the potential problems with this technique, three considerations relate to the assumption that R , the ox/red ratio, is constant. The diffusional mixing time within the cell must be fast relative to the kinetic run. Since mixing in a 0.15-mm cell occurs in ca. 20 s, this problem is avoided by assuring that the kinetic run lasts longer than a few minutes. Ohmic potential drop within the cell can lead to a
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
complexity of conventional electroanalytical techniques results from the simultaneous occurrence of diffusion, charge transfer, and kinetics. By removing diffusion from the picture, one obtains the simplicity of homogeneous kinetic methods combined with the additional advantage of potential control. It is anticipated that these important advantages will outweigh the potential problems alluded to above.
nonhomogeneous potential on the working electrode. This effect leads to deviations from the applied potential which are larger for large currents and would cause negative deviations from theory for higher R values. In this case, however, the current decreases greatly after the initial electrolysis, to the value required to maintain the oxlred ratio as the chemical reaction proceeds. This current is generally less than 10 PA, usually on the order of 1-3 PA, values which should result in negligible error in aqueous solutions. Further evidence that ohmic losses are negligible is provided by the fact that halving the PAP concentration, resulting in smaller total current, caused no observable change in rate constants at the half-wave potential. The third source of error is more difficult to avoid, that of a slow charge transfer rate. If the chemical reaction becomes fast relative to the charge transfer rate, the oxlred ratio will not be maintained, and the result will be a drawn-out kobsd vs. Eappplot, similar to a quasi-reversible polarogram. A likely means to remove this problem is the addition of a redox mediator to the solution. If the mediator interacts rapidly with both the electrode and the couple of interest, the solution potential will be maintained at the proper value, and the oxlred ratio will be a constant. Finally, it should be noted that all these effects can cause uncertainty in the starting time for the reaction. Although not important for first-order reactions, the uncertain starting time will be more critical for higher order systems. In these cases, care must be taken to assure that the kinetic run is long relative to the time required to establish the oxlred ratio. It should be emphasized that the strengths of this thin-layer kinetic method all result from the fact that one is working with a homogeneous solution in equilibrium with an electrode. The
ACKNOWLEDGMENT The author thanks William Heineman of the University of Cincinnati for helpful discussions during this work. LITERATURE CITED (1) R. N. Adams, "Electrochemistry at Solid Electrodes", Marcel Dekker, New York, N.Y., 1969. (2) A. C. Testa, and W. Reinmuth, Anal. Chem., 32, 1512 (1960). (3) W. N. Schwarz and I. Shain, J. Phys. Chem., 69, 30 (1965). ( 4 ) R. S. Nicholson and i. Shain, Anal. Chem., 36, 706 (1964). (5) M. Hawley and S. W. Feldberg, J. Phys. Chem., 70, 3459 (1966). (6) See, for example, L. Marcoux, J. Am. Chem. Soc., 93, 537 (1971). (7) T. Kuwana. Ber. Bunsenges. Phys. Chem., 77,858 (1973). (8) H. N. Blount, J. flectroanal. Chern., 42, 271 (1973). (9) R. W. Murray, W. R . Heineman, and G. W. O'Dom, Anal. Chem., 39, 1666 (1967). (IO) W. R. Heineman, 8.J. Norris, and J. F. Goetz, Anal. Chem., 47, 79 (1975). (1 1) Reference 1, p 336. (12) J. F. Corbett, J. Chem. SOC.5, 213 (1969). (13) L. Marcoux, J. Phys. Chem., 76, 3254 (1972). (14) C. R. Christensen and F. C. Anson, Anal. Chem., 36, 495 (1964).
RECEIVEDfor review September 7,1976. Accepted November 8,1976. The work was supported in part by a Cottrell Research Grant from the Research Corporation and by a grant from the Merck Company Foundation.
Activity Coefficients and Osmotic Coefficients in Aqueous Solutions of Choline Chloride at 25 OC J. B. Macaskill, M. S. Mohan,' and Roger G. Bates" Department of Chemistry, University of Florida, Gainesville, Fla. 326 1 1
Th? mean ionic activity coefficients of choline chloride (trimethyl-p-hydroxyethylammonium chloride) have been determined in aqueous solutions at molalities up to 4 mol kg-' by measurementswith an organic cation-selective electrode. The results are compared with those derived from a parallel isopiestic study of the osmotic coefflclents in these solutions at molalltles up to 7 mol kg-l. Agreement between the two methods is very good, indicating that the choline Ion-selective electrode shows near-perfect Nernstian behavior over the molality range 0.001 to 4 mol kg-I.
The choline ion is an essential component of biological systems. It is important in the metabolism of carbohydrates and nitrogen compounds and plays a role in lipid metabolism as well. The behavior of choline chloride (trimethyl-(3-hydroxyethylammonium chloride) has been studied previously by the gravimetric isopiestic technique. Fleming ( 1 ) concluded, on the basis of his experimental osmotic coefficient data, that the behavior of choline chloride in aqueous solutions Present address, D e p a r t m e n t of Chemistry, Texas versity, College Station, Texas.
A&M Uni-
is almost identical to that of ammonium chloride. This result is surprising in view of the structural similarity of the choline and tetramethylammonium ions and the established fact that the behavior of Me4NC1 in aqueous solutions is significantly different from that of NH&l(2). The osmotic coefficients (4) at a molality of 0.1 mol kg-l are 0.927 for NH4Cl and 0.910 for Me4NC1,and the difference increases with concentration. Boyd, Schwarz, and Lindenbaum ( 3 ) took note of this anomaly and made further isopiestic measurements on choline chloride. They concluded that choline chloride is indeed very similar to tetramethylammonium chloride and, in fact, that the osmotic and activity coefficients for choline chloride are even lower than those of Me4NC1. One might have expected any changes through replacement of a methyl group by a phydroxyethyl group not only to be slight but to be in the direction of increasing the osmotic and activity coefficients. Boyd et al. ( 3 )thus concluded that the (3-hydroxyethyl group must produce a sizable dipole moment in the cation. Whatever the explanation, the decreased activity of choline chloride relative to Me4NC1 is apparently substantiated by conductance studies ( 4 , 5 )which show that the limiting conductivity of the choline ion is significantly lower than that of MedN+. Baum (6) has made measurements on solutions of choline and acetylcholine chlorides using an organic cation-selective
ANALYTICAL CHEMISTRY, VOL. 49, NO. 2, FEBRUARY 1977
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