Heterogeneous mechanisms of the oxidation of catechols and

Mohan Kumar , B. E. Kumara Swamy , Sathish Reddy , T. V. Sathisha , J. Manjanna. Anal. Methods 2013 5 (3), ... Analytica Chimica Acta 2012 721, 55-60 ...
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Anal. Chem. 1966, 58,1474-1480

dominating peak a t 6.3. The affinity spectrum also indicates that there may be a functional group with a pK value around 8, which is not seen in the integral pK spectrum. Figure 5, again, illustrates the rather large differences between smoothes and unsmoothed data. The peak at a pK value of 6 is clearly visible in the spectrum of the smoothed data, while it is much less pronounced for the unsmoothed data. In concluding, we believe that the integral pK spectrometry provides an interesting and useful alternative to the so far published methods. The main disadvantage of our method is the linear least-squares regression with bounding constraints that requires somewhat elaborate calculations. The main advantage of our method is certainly the clearly defined resolution in pK units.

ACKNOWLEDGMENT The sample of fulvic acid studied in this work was kindly supplied by H. H h i , Institute for Agricultural Research, Bern, Switzerland.

LITERATURE CITED Arp, P. A. Can. J . Chem. 1983, 6 1 , 1671-1682. Eberle, S. H.; Feuerstein, W. Naturwissenschaften 1979, 66, 572-573. Gamble, D. S. Can. J . Chem. 1970, 4 8 , 2662-2669. Arena, G.; Rlzzarelli, E; Sammartano, S.; Rigano, C. Talanta 1979, 26, 1-14. Fianagan, M. T.; Tattam, F. G.; Green, N. M. Immunochemistry 1978, 15, 261-267. Shuman, M. S.; Colllns, B. J.; Fltzgerald, P. J.; Olson, D. L. In Aquatic and TerrestrialHumlc Materlals; Chrlstman, R. F., Gjessing, E. T., Ed.; Ann Arbor Science: Stoneham, MA, 1963; pp 349-370. Thakur, A. K.; Munson, P. J.; Hunston, D. L.; Rodbard, D. Anal. Biochem. 1980, 103, 240-254. Hunston, D. L. Anal. Biochem. 1975, 63, 99-109. Ferry, J. D. Viscoelastic Properties of Polymers; Wiley: New York, 1961; p 63ff. Forsling, W.; Hietanen, S.; Sillen, L. G. Acta Chem. Scand. 1952, 6 , 901-909.

Received for review July 11, 1985. Resubmitted December 16, 1985. Accepted January 23, 1986. The work was financially supported by the Swiss National Foundation for Scientific Research.

Heterogeneous Mechanisms of the Oxidation of Catechols and Ascorbic Acid at Carbon Electrodes Mark R. Deakin, Paul M. Kovach, K. J. Stutts, and R. Mark Wightman* Department of Chemistry, Indiana University, Bloomington, Indiana 47405

Evaluatlon of the pathways for the oxldatlon of organlc substances has been used to probe the differences between actlvated and unactlvated carbon electrodes. The oxldatlon of some substituted catechols and ascorbate has been examlned at lntermedlate pH. The heterogeneous rate constants for the reactions have been evaluated by use of semllntegratlon. Consideration of the pK,’s of the lntermedlates and the formal potentlals for the one-electron processes has enabled the microscopic pathways to be determlned. At pH 7.0, catechol oxldatlon proceeds through the one-electron oxldatlon of the monoanlon followed by oxldatlon of the radlcal anion. Identlcal mechanlstlc behavlor is found for the oxldatlon of ascorbate. The rates for the one-electron processes are slower at unactlvated carbon electrodes than those observed at metal electrodes and heat-treated carbon, or as calculated from the homogeneous electron transfer rates. Thus, It Is concluded that surface actlvation does not result In electrocatalysis, but rather removes the impedlments to electron transfer that exlst at unactlvated electrodes.

Numerous observations have been made concerning the effect of various electrode pretreatments on the electrochemical response observed a t carbon electrodes ( I ) . These procedures are of interest to evaluate different mechanisms of electron transfer and also to try to obtain electrocatalysis a t carbon electrodes. Although there have been numerous elegant attempts to modify carbon surfaces, it has been found that improvements in voltammetric behavior can be obtained by simple pretreatment of the electrode surface. Pretreatments that have been useful include high temperatures @), electrochemical anodization (3-8), and high-speed polishing methods (9-10). These treatments are efficacious at several

forms of carbon such as glassy carbon, carbon paste, and carbon fibers (1). This is not surprising since, although the interior structure of different forms of carbon may differ, heterogeneous electron transfer would be expected to be most affected by the exterior of the material. The surface analysis of carbonaceous materials demonstrates that a large amount of covalently bound oxygen is present at most exposed interfaces (4, 6, 10-13). Various mechanisms have been proposed to explain acceleration of heterogeneous electron transfer rates observed after surface pretreatment of carbon electrodes. All of these treatments appear to change the structure and number of oxygen-containing functional groups on the electrode surface (3, 4, 10). Functional groups have been proposed to affect electron transfer either by promoting hydrogen atom transfer (6,10) by altering the charge at the electrode surface (10, 14) or by providing redox sites at the electrode surface that could be involved in mediation of electron transfer (8, 11-13). It also has been observed in all of these studies that the apparent capacitance of the electrode changes, perhaps indicating an increase in the microscopic area of the electrode surface (1-4). An alternate interpretation is that heterogeneous electron transfer rates are impeded by the hydrophobic nature of the surface (5, 15) and that this hydrophobicity is altered by surface pretreatments (3, 4). Despite the fact that many mechanisms have been proposed for the activation seen after certain treatments a t carbon electrodes, there has been very little work to understand the microscopic kinetic steps that occur at these electrodes. Through analysis of the kinetics of highly charged outer-sphere reactants, we have shown that the differences between carbon surfaces and metallic electrodes can be explained in part by the presence of a small number of carboxyl groups on untreated carbon surfaces (14). These can be removed by subsequent treatments. It has also been shown through kinetic

0003-2700/86/0358-1474$01.50/00 1986 American Chemical Society

ANALYTICAL CHEMISTRY, VOL. 58, NO. 7, JUNE 1986

isotope effects that protons can play an intimate role in some heterogenous electron transfer reactions (6). However, rigorous kinetic analyses of proton-coupled redox reactions a t carbon surfaces have not been reported. Here, we present data concerning the oxidation of some catechols and ascorbic acid at carbon electrodes and at neutral pH. These reactions are two-electron oxidations that also involve proton transfer. The primary tool employed in this work is semiintegral analysis. This method of kinetic analysis allows the rate of electron transfer to be measured over a range of potentials near the formal potential of the redox couple. This rate information can be used two ways. First, for the two electron couples examined here, the shift from second-electron ratelimiting behavior to a first-electron rate-limiting situation can be identified. This shift should occur at high overpotentials for all the two-electron couples examined (16). Second, when the first electron is rate limiting, variation of solution pH can be used to test whether or not proton transfer precedes electron transfer. The rate constants and mechanisms obtained in this manner a t conventional carbon surfaces can be compared to those obtained at activated carbon surfaces, at metal electrodes, or in homogeneous solution to further define activation mechanisms at carbon.

EXPERIMENTAL SECTION Chemicals and Solutions. Experimentswere performed with water purified by ion exchange and distillation in a Corning Mega-Pure water purification system. Ascorbic acid, 3,4-dihydroxyphenylaceticacid (DOPAC),3,4-dihydroxybenzylamine (DHBA), and 2,4,5-trihydroxyphenethylamine(6-hydroxydopamine, 6-OHDA) (Sigma,St. Louis, MO) were used as purchased. The 4-methylcatechol(4-MC) (Aldrich, Milwaukee, WI) was recrystallized from hot toluene. The concentration of the organic M in buffered solutions. Low concencompounds was 2 X trations were used to minimize filming effects. Buffers were prepared from 0.1 M solutions of mono- or dibasic sodium phosphate adjusted to the desired pH with phosphoric acid. The ionic strength of the buffers was adjusted to 1.0 with sodium chloride. Solution pH measurements were made preceding and following voltammetricexperiments with a Corning Model 12 pH meter. Measurements were restricted to pH values less than 7.5 because of the instability of these compounds with respect to molecular oxygen and hydroxide. Electrodes and Apparatus. Glassy carbon electrodes were prepared as described earlier (2).“Unactivated” electrodes were hand polished until smooth, and before each run polished with 0.05-~malumina and refluxed in toluene or methylene chloride. Carbon paste electrodes (radius of 0.8 mm) employed a paste containing 66% graphite powder (Ultra Carbon, Bay City, MI) and 34% silicone grease (Dow Corning, Midland, MI) that were thoroughlymixed. Carbon paste was employed for kinetic analysis because it exhibits a low charging current (5), which facilitates semiintegral analysis. The electrodes were resurfaced after each voltammetric scan. A hanging mercury drop with an area of 4 X cm2 also was used. Digital acquisition of data was accomplished with a system previously described (14).A saturated sodium calomel reference electrode and a platinum wire auxiliary were employed, and experiments were conducted at 23 k 2 O C . All potentials are reported vs. the sodium saturated calomel electrode (SSCE). Voltammetric Analysis. Due to slight variations in the preparation of the carbon paste electrodes, average cyclic voltammograms were used in the kinetic analysis. Ten cyclic scans, at 0.1 V s-’, were obtained and combined for each value of solution pH. Each scan was recorded at a fresh electrode surface. Semiintegrationof the averaged cyclic voltammograms followed the method of Oldham (17). Current measurements, made at 1.2-mV intervals, were integrated with respect to the inverse square root of time using the “G1 algorithm”. Analysis of the semiintegral to obtain the rate constants followed the method of Imbeaux and Saveant (It?), and a value of n = 2 has been applied in accord with the known chemistry of catechols and ascorbate (29). The rate of electron transfer &(E))for quasi-reversible systems on the positive sweep may be expressed as a function of

DHBA

I

I

I

0.6

I

I

0.3 0.0 -0.3

,

V

I

OOPAC

.

,

1475

I

,

0.6 0.3 0.0 -0.3 E (V vs SCE)

I

0.6

,

,

0.3 0.0 -0.3

Figure 1. Cyclic voltammograms of 4-MC, DHBA, and DOPAC at 0.1 V s-’ in buffered solution at intermediate pH: (top)at carbon paste electrodes with semiintegrated current (I)and (below)at glassy carbon electrodes. the semiintegral (I)and the current ( i ) at different values of the applied potential (E) k(E) = D1IZ i (1) Ilim - Z(E)(1+ exp[nF/RT(E’’ - E ) ] ] In this expression, D is the diffusion coefficient and Eo’is the formal potential of the two-electron process. On the reverse sweep ( E O ’ - E ) is replaced by (E - EO’), and the zero point for Z(E) is reset to the start of the reverse wave. For electrochemically irreversible systems the relationship for the forward wave is Diffusion coefficients were taken from ref 20. These equations require the assumption that the radical intermediate only react with the electrode to form products (vida infra). For mechanisticanalysis a Volmerian model has been assumed. The Volmerian assumption is that the activation energy of the reaction is equally partitioned between products and reactants. Thus, for a two-electron process the rate-determining step can be deduced by examining the slope of In k ( E ) vs. E. If the first electron transfer is rate determining the slope is 0.5F/RT, while if the second electron transfer is rate determining the slope is 1,5F/RT. The experimentaluncertainty in the rate measurements has been estimated by calculating the values of In k ( E ) for each of the 10 cyclics used to give the averaged cyclic. The standard deviation in In k(E) by this process is 0.2. The formal potential of the two-electron process (Eo()can be evaluated for quasi-reversible systems at each pH from the cyclic voltammetric and semiintegral data. On the return sweep of the cyclic voltammogram, i passes through zero. At this point the formal potential of the couple may be calculated (21)with the following expression: Eo’ = Ei-0 - (RT/nF) In

[(Zlim

- lt=O)/lL=O1

(3)

RESULTS Voltammetry of Catechols. At neutral values of pH the shape of voltammograms of catechols obtained at unmodified carbon surfaces is indicative of an electrochemically quasireversible process. Voltammograms for three catechols, which differ only in the charge of the side chain, are given in Figure 1, along with the semiintegrated form of these data. Voltammograms appear similar for all of the catechols at this formulation of carbon paste. In contrast, voltammograms for DOPAC are more irreversible than the other two catechols at unactivated glassy carbon electrodes. The shape of the cyclic voltammograms is indicative of the microscopic reactions that are occurring during electrolysis. The kinetics of these reactions were evaluated with semiintegral analysis. Inspection of the semiintegrals obtained a t carbon paste electrodes indicates that the cathodic and anodic

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Table I. Formal Potentials for the Two-Electron and One-Electron Processes for Various Catechols and the Measured Rate Constants at the Two-Electron Formal Potentials’ compd

PH

4-MC

6.03 6.9 6.40 7.50 6.61 7.60

DHBA DOPAC

0.159 0.108 0.211 0.146 0.135 0.076

-0.045 -0.046 0.048 0.048 -0.029 -0.030

0.363 0.262 0.374 0.244 0.299 0.182

4.7 2.6 2.3 0.94 2.6 0.90

2.8 1.3 1.3 0.90 1.2 0.30

‘Rates are measured at carbon paste electrodes. bCalculatedfrom data in ref 23 and 24 and Table I1 using E0’2,app = Eo + (RT/F) In (Ka~/Ks6)[([H+I/Kad -k 1 -k Ka6/[H+I)/([H+I/Kaz+ 1 + &4/[H+I)I and E”’1,spp = 2E”’ - E”’2,app.

2

based on consideration of the relevant pKa)s, the initial and final products are those shown in reaction 4 for the conditions used in this work. The one-electron pH-dependent rate constants (ko’L,app) can be evaluated by using the following expressions (16). For the reduction,

0

- 1- -

EX:W p H 6.0

pH 6.9

EP,’a,, pH 6.0,6.9

1

I

k(E)

k I” D1/2 - 2 -4

1 kO’l,app exp(-O.5F/RT(E - Eo’l,app)) 1 (5) k*2 exp(-1.5F/RT(E -Eo’))

+

k*2 = k0’2,app exp(-0.5~/RT(Eo’2,app - EO’))

(6)

while for the oxidation,

-6

0.4

0.3

0.2 0.I E ( V vs SCE)

0.0

-0.1 1

Figure 2. Rate of heterogeneous electron transfer for 4-MC at a carbon paste electrode in pH 6.0 (0,A)and pH 6.9 (0, A) buffered anodic data, (A,A) cathodic data, (-) calculated SOlUtiOn: (0,0) curve from eq 5-8 and values given in Table I;scan rate 0.1 V s-’, 10 repetitions each.

sweeps are similar in amplitude at values of pH near 7 . This clearly shows that the redox processes are not distorted by accompanying dimerization or other chemical processes. No significant surface waves were observed; thus, adsorbed species are not major contributors to the current. The two-electron formal potentials (EO’) evaluated from the semiintegrals are given in Table I. The values for 4-MC are in good agreement with those reported previously (22).

Rate Constants for the Apparent One-Electron Processes. The semiintegral method is unique in that the rate constants for both processes can be observed even though the bulk solution only contains the reduced form (QH,) of the couple. Evaluation of the rate constants from the semiintegral data for the forward and reverse wave (eq 1) leads to curves similar to those shown in Figure 2. The anodic and cathodic branches intersect at the formal potential for the two-electron couple ( E O ’ ) . Data for all three catechols show a nonlinear dependence of k ( E ) on potential. This is consistent with a heterogeneous electron transfer reaction in which the ratedetermining step changes as a function of potential (16). For the catechols, which are known to be oxidized via a twoelectron process, the reaction sequence can be written as (4) where R denotes a radical intermediate and Eo’1app represents the formal potential for the one-electron processes at a specific pH. The formal potentials can be calculated from literature values (Table I). This scheme gives no indication of the order of electron and proton transfer or the nature of R. These will be determined by evaluation of the kinetic data. However,

k*l exp(-1.5F/RT(Eo’ k*i =

kO’l,app

- E)) (7)

exp(-0.5F/RT(Eo’ - Eo’l,app)) (8)

At low overpotentials the final term in eq 5 and 7 predominates. Therefore, the second electron transfer is rate determining, and a plot of In k ( E ) vs. E has a slope of 1.5FIRT. Alternately, at high overpotentials the initial term on the right-hand side predominates, and the first electron transfer becomes rate limiting. For these compounds the anodic and cathodic branches of the rate plots at the values of Eo’l,app and E0‘2,app have a slope of 0.5, indicating first-electron rate-limiting behavior at these potentials (Figure 2). Thus, the values of k(E) yield directly values for ho’l,appand k0’2,app in scheme 4 (Table I). Extrapolation of the rate constants to other potentials is done with eq 5-8. The results of this analysis are the solid curves given in Figure 2. Evaluation of the Reaction Sequence. The relative order of the proton and electron transfers for this system can be inferred from the pH dependence of the rate data discussed above. Inspection of the anodic data in Figure 2 shows that the curves are vertically displaced by 1.5 log units at high overpotentials. A shift of 2.0 log units would be expected for a first-order dependence on protons for this pH change. A t high overpotentials on the cathodic branch a smaller change in the vertical position of the rate constant curves (-0.5) is observed, implying a near-zero-order dependence of this reaction on pH. Thus, although an exact mechanism cannot be discerned, these observations provide a preliminary indication that the first-electron oxidation of the catechols is first order in protons, but the first-electron reduction of the oquinone is zero order in protons. A more complete kinetic analysis of data can be obtained by using the method published by Laviron (25-27). The analysis requires consideration of all of the possible intermediates, the pK:s of the relevant products and intermediates, and the EO’S for the single electron transfers. The overall

ANALYTICAL CHEMISTRY, VOL. 58, NO. 7, JUNE 1986

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Table 11. pK, Values for the Systems Examined” P K ~ I PK~, PKaa

QH, 4-MC

-10

-1

-6

DHBA DOPAC

-10

-1 -1

-6

-10

-6

AA

PKa,

PKa5

PKa6

5’ 4.7d 5‘ -0.4e

9.6‘

12.5‘ 12.5d 12‘ 11.57f

8.gd 10‘ 4.17f

”Unlessnoted the values are typical for stable quinones (23-25). Reference 28. Reference 29. Values for 4-(aminoethyl)catechol, ref 30 and 31. eReference 32. fReference33. Table 111. Microscopic Heterogeneous Rate Constants for Three Catechols at Carbon Paste Electrodes compd

PH

k3, cm 5-l

k5, cm e’

4-MC

6.03 6.90

0.046

6.6

DHBA

6.40

0.028 0.032

DOPAC

7.50 6.61 7.60

2.75 2.0 1.2 4.9 1.05

0.014 0.028 0.010

reaction scheme can be represented in a “nine-member box”. The notation in this scheme follows that of Laviron. Values of the pK,’s for many of these intermediates are available (Table 11). When values were not available, those for similar compounds have been employed. The pK,’s of the radicals have been determined by pulse radiolysis experiments (28, 30,31). It can be shown that the extreme values of the pK,‘s for QH$+ and Q2-make them unlikely intermediates at neutral pH values (25). However, there are still three possible radical intermediates. The importance of each can be evaluated from the kinetic data and the estimated pK,’s.

0.0

0.6

0.4

E

0.2

is not expected. The slight decrease in the computed microscopic rate constants as a function of pH may arise from an underestimationof the importance of k2 or kg. For example, evaluation of the importance of kz depends in part on the accuracy of the value for Ka3,a value which has been estimated. However, if a first-order dependence on protons exists that has not been accounted for, the rate constants would vary by a factor of 10 per unit of pH change. As will be shown elsewhere, these rate constants are consistent with values obtained over the pH range of 1-7.5 (34). The observed values of the rate constants also are dependent on the type of carbon paste employed (5). Since the computed values for k3 and k5 are relatively pH independent for all three catechols, the same pathway must be followed by all. The data at neutral pH are consistent with the following pathway: -H+

kl

E O 1

’

Q H ~ + *’

k4 EO4

’

QH2

The rate constants given in Table I are a composite of the primary processes noted above

ko’l,app =

(kz

(1

+ k3(Ka$a4)1’2[H+l-1)

+ Ka3[H+]-1)1/2([H+]Ka2-1 + 1 + Ka4[H+]-1)1/2(10)

k0’2,app= (k,(Ka,Ka5)-1’2[H+I+ k5) ([H+]Ka2-l 1 + Ka4[H+]-1)1/2([Hf]Ka6-1 + 1)1/2(11)

+

where the k:s are the microscopic rate constants as given in the “nine-memberbox” scheme (25). Voltammograms for each catechol were obtained at two different pH values, and thus, two values of each koi,appwere available to evaluate the microscopic rate constants (Table 111). A t the pH values employed, the contribution from the pathways given by k , and k4 is sufficiently small that only k3 and k5 can be calculated. The evaluation of the microscopic, pH-independent rate constant allows a direct comparison to be made with other one-electron systems (vide infra). The calculated values in Table I11 are directly dependent on the uncertainties of the pK,’s and the computed Ei0”s. Thus, while the rate constants are self-consistent, precision

-0.2

Flgure 3. (Above) Voltammetry of AA at a carbon paste electrode in pH 7.0 buffered solution at 0.1 V s-’. (Below) Rate of heterogeneous electron transfer from semiintegral analysis.

-e

-H+

QH2 =i QH- =i QH. % Q-.

Q H ~ ~ +

00

( V vs S C E )

Q

This scheme is dominant for the benzoquinone/hydroquinone couple at platinum electrodes (25, 35). The chemistry of these redox systems is now defined at a carbon paste surface. Cyclic voltammograms of these same catechols at glassy carbon are similar in shape, and thus, an indication is given that the same reaction scheme is followed. The slope of the cathodic wave for DOPAC is steeper at glassy carbon than at carbon paste. This is indicative that k5 is rate limiting for a greater potential range (Le., k3 is faster than at carbon paste). Oxidation of Ascorbic Acid at Carbon Electrodes. Ascorbic acid is also oxidized in a two-electron process, but the oxidation of ascorbate is accompanied by a rapid hydration reaction (19). Thus, the cathodic wave for this redox process is only observed at very fast sweep rates. At carbon paste electrodes and polished glassy carbon electrodes (Figure 3) the voltammetric wave is extremely irreversible. Voltammetric data from carbon paste and glassy carbon electrodes have been semiintegrated,and the limiting semiintegralobserved is close to that expected for n = 2. The slope of the data for In k ( E ) vs. potential is linear at both surfaces with a slope near 0.5 (carbon paste, slope = 0.436, correlation coefficient = 0.9995, pH 6.72; glassy carbon, slope = 0.352, correlation coefficient = 0.982, pH 7.0). Kinetic analysis of the type used for the catechols is not possible for the oxidation of ascorbate since the reverse wave is absent. However, if a rapid chemical reaction follows heterogeneous electron transfer, the position and shape of an electrochemical wave can be solely dependent on the elec-

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0.4

0.3

I

r

A A

4

A

A

5

6 PH

7

8

Flgure 4. Oxidation of AA at a carbon paste electrode: (above)

transfer coefficient and (below) peak potential.

trochemical kinetics (36,37). This is true when the ratio k-,/k, approaches zero, where k-, is the rate constant of the reductive electrochemical process at the peak potential (E,) and k , is the sum of the rate constant k , for chemical decomposition and kd the rate constant for diffusion. The value for kd is given by kd = 2.82(nDu)lI2where u is the scan rate. The rate of decomposition, k,, can be evaluated from k , = (Dkf)lIzwhere kf is the rate of the hydration reaction (1.4X lo3s-l (38,39)). For the oxidation of ascorbate k, > 30kd,which supports the assignment of an electrochemicallyirreversible process at the peak potential, E,,. The rate k1is evaluated by first estimating the apparent standard rate constant for the reaction. The value of the apparent electrochemical rate constant on the anodic branch at 0.0 V can be determined from the semiintegral data. At a pH of 6.72and at carbon paste electrodes, k(0.0 V) is equal to 1.2 X cm s-l. Extrapolation of k ( E ) with the Volmerian assumption can be used to estimate a rate constant at either the first-electron potential ( E O ’ = 0.059V, pH 7 (23,24))or the second-electron potential ( E O ’ = -0.182 V, pH 7 (40)). Evaluation of the reverse rate constant at the peak potential of the cyclic voltammogram (EP= 0.36V, pH 6.7)from either of the calculated standard potentials gives a value sufficiently less than k , so that the ratio k q / k , is effectively zero. Consequently, for oxidation of ascorbate at carbon paste and glassy carbon electrodes where the first electron transfer is rate determining, E, is indicative of the events accompanying that transfer. The change in this parameter (Up) with pH is given by (41)

AE,=-

RT

0.5F

A In kapp + M i o

Here, Eio is the appropriate one-electron standard potential. If proton transfer occurs after the first electron oxidation, then Eo5of scheme 9 is this value, and AE, should be 0 for changes in pH under conditions where the reaction mechanism and rate remains unchanged. The peak potential was examined as a function of pH from pH 4.75 to 7.75 at carbon paste electrodes to determine if this is the case. The average value for the peak potential was 0.345 V (Figure 4),with values at all pHs clustered around the mean. Semiintegration and examination of In k(E) as a function of potential at each pH gives a slope near 0.5. These data indicate the microscopic pathway for the oxidation of ascorbate at carbon electrodes. At neutral pH values the primary form of ascorbate is the monoanion (AH-). The first-electron oxidation of this species can be assigned to be rate determiningat carbon paste and unmodified glassy carbon (if a Volmerian model is assumed) because of the dependence

0.6

0.4 E

0.2

0.0

-0.2

( V v s SCE)

Figure 5. (Above)Voltammetry of AA (1 mM) at a heat-treated glassy carbon electrode in pH 7.0 buffered solution at 0.1 V s-’. (Below)Rate

of heterogeneous electron transfer from semiintegral analysis. of k ( E ) on EaPvFurthermore, the oxidation is independent of pH. Therefore, the proposed pathway is -H+

AH- 5 AH. =, A-.

5A

The electrochemical data only give information up to the generation of AH.. However, the route through the radical anion (A-.) is dictated in light of the pKa of the neutral radical (Table 11). Furthermore, this species has been observed in pulse radiolysis experiments at neutral pH (42). After formation, the radical anion undergoes a one-electron oxidation and hydrolysis (19).

Oxidation of Catechols and Ascorbate at Activated Carbon Surfaces.. Esterification of the surface of glassy carbon has been shown to reduce barriers for heterogeneous electron transfer of negatively charged, outer-sphere complexes. This treatment caused a small change in the peak potential for ascorbate (EP = 0.161 V at esterified glassy carbon; E, = 0.193 V at polished surfaces, both at 100 mV s-,) corresponding to an increase in the rate of 46%. In contrast, this treatment causes at least a 100% increase in the rate constant for the reduction of ferricyanide. Chemical reduction of glassy carbon also has been shown to increase the degree of reversibility of ferricyanide at neutral pH at glassy carbon surfaces. No effect on the position of E, for ascorbate was noted for this treatment within experimental error. In contrast, heat treatment of a glassy carbon electrode results in sharper voltammetric waves for all of the compounds described in this work. At moderate scan rates, kinetic information cannot be obtained for the catechols since they are almost reversible (Up = 50 mV). However, semiintegral analysis of the voltammograms for ascorbate (Figure 5) does yield useful kinetic information. The slope of the measured In [ k ( E ) ]vs. Ea,, has a slope of 1.36 at the foot of the voltammetric wave that changes to 0.5 at higher overpotential. These data indicate that the rate of the first electron transfer has been accelerated compared to an unmodified electrode. Thus, the second electron transfer becomes rate determining at low values of overpotential, and a t higher values a oneelectron process is rate determining. Voltammograms of organic compounds at heat-treated electrodes tend to resemble those at clean metal electrodes. The position of Epfor ascorbate is identical with that observed at mercury electrodes ( 1 ) . Eo’values for catechols at intermediate pH are too positive to be examined with mercury electrodes. However, 2,4,5-trihydroxyphenethylaminehas been examined (Figure 6). The voltammograms for this

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homogeneous electron exchange rate constant (kex)has been measured to be 6 X lo’ M-l s-l in aqueous solution at pH 7 (43). Catechols follow similar kinetic behavior (23,24). The value for the benzoquinone couple can be used to calculate an approximate value for the heterogeneousrate constant, k,, using the Hush relationship (ignoring double-layer effects) (46, 48) kg = (3 x 10-8) k,, (15)

+0.2

-a2

-a6

+02 -0.2 E ( V K SCE)

-06

-a2

-a6

Figure 8. Voltammetry of 6-OHDA at various electrodes in pH 7.4 phosphate-citrate buffer at 0.1 V s-’: (a) untreated glass carbon, (b) heat-treated glassy carbon, and (c) mercury.

compound a t unmodified glassy carbon are similar in shape to those of the catechols, suggesting a similar reaction scheme. Heat treatment of the electrode increases the reversibility of this compound and provides an electrode with electrochemical properties more similar to those observed a t mercury electrodes.

DISCUSSION Semiintegral analysis of voltammograms for catechols a t carbon paste provides sufficient information to identify the intermediates involved in the heterogeneous reaction. The mechanism identified at neutral pH for the oxidation involves the intermediacy of QH-, QH., and Q-e. The order of proton and electron transfers is in agreement for the benzoquinone/ hydroquinone couple at a platinum electrode (25, 35), and the mechanism is also in accord with that found in a homogeneous environment (43). The validity of the analysis used to determine this mechanism rests on four assumptions: that a Volmerian model for charge transfer is applicable, that current from adsorbed species is negligible, that the radicals do not disproportionate, and that the accompanying proton transfers are rapid. The latter assumption appears valid, since proton transfers are known to occur under diffusion control in aqueous solutions (44). The experimental data support the Volmerian assumption. In addition, a large number of oneelectron processes for the generation of organic radicals have been shown to have a transfer coefficient of -0.5 (45). The well-defined nature of the semiintegral indicates that diffusion-controlled electrochemistry is the predominant contributor to the current. The disproportionation of radicals under similar conditions has been evaluated and found negligible (26). The three catechols were selected for study because of the difference in charge of their side chains at intermediate pH. Inspection of the microscopic rate constants (Table 111) indicates that there is not a significant difference for these compounds with the carbon paste employed. This indicates that the interaction of surface carboxyl groups, with the side chain portion of the catechol, is minimal. In contrast, the shape of the voltammograms at glassy carbon electrodes indicates a slower rate for DOPAC than the other catechols, which implies an effect of surface anionic groups on the heterogeneous electron transfer. This observation has also been made with carbon fibers and carbon paste with Nujol as binder (45). The microscopic heterogeneous rate constants for the first reduction step of all the o-quinones at carbon paste are less than 0.05 cm s-’. In contrast, the measured heterogeneous rate constants for the one-electron reduction of p-quinones in aprotic solvents is 10-50 times larger at metal electrodes (46, 47). For the benzoquinone/radical-anion couple, the

The value calculated in this manner is also 50 times greater than those measured at carbon paste electrodes. Thus, the heterogeneous rate constants for all three catechols are lower at carbon paste than would be expected at an ideal electrode. One interpretation of the reduced rate constant at unactivated electrodes is that the surface of the electrode is partially blocked (49, 50). At a partially blocked surface comprised of closely spaced sites, rate constants appear to be retarded because of the reduced area available for electron transfer. A physical interpretation of the origin of this effect is possible at carbon paste electrodes since the carbon particles are imbedded in an inert binder ( 5 ) . The voltammetric evidence also shows that the same “blocked” appearance is present at unactivated glassy carbon. The rate constants reported here for the one-electron processes are comparable in magnitude to those calculated by Laviron for the benzoquinone/hydroquinone couple at platinum electrodes (25). However, these measurements were probably made at a blocked platinum surface as hydroquinone has been shown to strongly adsorb at platinum surfaces (51-53). If the adsorption of anions is blocked by prior adsorption of iodine atoms, much faster rates are observed (54). The semiintegral technique also has been useful for the evaluation of the reaction mechanism of ascorbate. The same assumptions were made for this analysis as for the catechols. The experimentaltechnique does not permit evaluation of Eo’ for this case because the oxidation is followed by a chemical reaction that is very fast on the time scale of the electrochemical experiments. However, kinetic information can be obtained that permits the evaluation of the rate-limiting reaction step. A t unactivated carbon electrodes the oxidation of ascorbate to the free radical is rate limiting. At heat-treated glassy carbon the rate-limiting step for the oxidation shifts to the second-electron transfer in accord with the behavior observed at metal electrodes (55-58). At higher values of overpotential, the rate-limiting step for oxidation of ascorbate reverts to a fist-electron rate-limiting step. The change in rate for the first-electron transfer for the oxidation of ascorbate at heat-treated carbon surfaces can be estimated with eq 13. Since activation by heat treatment of the electrode causes -300-mV shift in the peak potential, the rate constant for the first-electron oxidation must have been accelerated 300-fold. This is a lower limit, since the firstelectron transfer is no longer rate limiting at the activated surface. If activation by heat treatment results in an increase in microscopic area, comparable enhancements should be observed for other compounds (50,51). However, this is not the case for reduction of Fe(CN)63-where the rate is only enhanced 2-4-fold by heat treatment (10). The change in rate for at some treated carbon surfaces (prepared by chemical reduction or surface esterification) has been attributed to electrostatic effects (14), and similar causes may also affect the slight rate enhancement for ascorbate at these surfaces. However, the differences in acceleration of rate for these two compounds at heat-treated electrodes are too large to be explained by electrostatics. Thus, it appears that some property other than microscopic area changes or electrostatics plays a role in the observed effects. One proposed model for electrode activation is that charge transfer involves hydrogen atom transfer ( 6 , 9 ) . Evidence for

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NO. 7, JUNE 1986

this model comes from double waves seen for catechol oxidation in acidic media, which are attributed to radical intermediates. However, this mechanism is implausible, since the equilibrium constant for radical disproportionation is lo9 (28) and the rate of disproportionation of quinone radicals is several orders of magnitude greater than can be observed electrochemically (23,24,28). Double waves can arise with electrodes that are only partially activated as has been observed following pulsed laser treatment of surfaces (59) or proposed after electrochemical activation (60, 61). Since many surface treatments are efficacious for the activation of carbon surfaces for voltammetry, it is unlikely that there exists a single mechanism that explains all. However, ascorbate, the most polar compound examined in this work, exhibits the largest acceleration, which supports the proposal that the hydrophobic nature of unactivated surfaces is a major contribution to the observed results. Furthermore, electrode activation does not affect the rate of nonpolar substances (60, 61) and has only a small effect on couples in nonaqueous solvents (10). The deterioration of activated heat-treated surfaces may result from a “reblocking” of hydrophilic sites on the electrode surface (62,63). To provide electrodes useful for electroanalysis, methods must be developed to prevent this deterioration. Registry No. 4-MC, 452-86-8; DHBA, 37491-68-2;DOPAC, 102-32-9;6-OHDA, 1199-18-4;AA, 50-81-7;Hg, 7782-42-5;carbon, 7440-44-0;graphite, 7782-42-5. LITERATURE CITED Wightman, R. M.; Deakin, M. R.; Kovach, P. M.; Kuhr, W. G.; Stutts, K. J. J. Nectrochem. SOC.1984, 131, 1578-1583. Stutts, K. J.; Kovach, P. M.; Kuhr, W. G.; Wightman, R. M. Anal. Chem. 1983, 55, 1632-1634. Engstrom, R. C.; Strasser, V. A. Anal. Chem. 1984, 56, 136-141. Engstrom, R . C. Anal. Chem. 1982, 54, 2310-2314. Rice, M . E.; Galus, 2 . ; Adams, R. N. J. Nectroanal. Chem. 1983, 143, 89-102. Cabaniss, G. E.; Diamantis, A. A.; Murphy, W. R,, Jr.; Linton, R. W.; Meyer, 1.J. J . Am. Chem. SOC.1985, 107, 1845-1853. Gonon, F. G.; Fombarlet, C. M.; Buda, M. J.; Pujol, J. F. Anal. Chem. 1981, 53, 1386-1389. Cenas, N.; Rozgaite, J.; Pocius, A,; Kuiys, J. J . Electroanal. Chem. 1983, 154, 121-128. Thornton, D. C.; Corby, K. T.; Spendel, V. A.; Jordan, J.; Robbat, A,, Jr.; Rutstrom, D. J.; Gross, M.; Ritzier, G. Anal. Chem. 1985, 57, 150-155. Kamau, G. N.; Wiiiis, W. S.; Rusling, J. F. Anal. Chem. 1985, 5 7 , 545-55 1. Evans, J. F.; Kuwana, T.; Henne, M. T.; Royer, G. P. J . Electroanal. Chem. 1977, 80, 409-416. Evans, J. F.; Kuwana, T. Anal. Chem. 1977, 49, 1632-1635. Miller, C. W.; Karweik, D. H.; Kuwana, T. Anal. Chem. 1981, 53, 2319-2323. Deakin, M. R.; Stutts, K. J.; Wightman, R. M. J. Nectroanal. Chem. 1985, 182, 113-122. Morcos, I.J. Chem. Phys. 1972, 5 7 , 1801-1802. Albery, J. Electrode Kinetics; Clarendon: Oxford, 1975; Chapter 5. Oldham, K. B. J. Electroanal. Chem. 1981, 121, 341-342. Imbeaux, J. C.; Saveant, J. M. J. Nectroanal. Chem. 1973, 44, 169- 187. Dryhurst, G.; Kadish, K. M.; Scheiler, F.; Renneberg, R. Biological Nectrochemistry; Academic Press: New York, 1982; Vol. 1. Gerhardt, G.; Adams, R. N. Anal. Chem. 1982, 54, 2618-2620.

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RECEIVED for review October 17, 1985. Accepted February 3,1986. This research was supported by NIH (R01 NS 15841).