Anal. Chem. lS86, 58,3235-3239
3235
I! 1.0
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Figure 3. Effects of solvent composition (chlwoform-heptanol) on the relative extraction efficiency of Lu3+/La3+with symdibenzo-16-
crownd-oxyacetic acid. cations. The total concentrations of Lu3+and the competing lanthanide were kept higher than that of the chelating agent so that only partial extraction of the lanthanides would take place. The peak of 177Lu(208 keV) is quite different from those of the other rare-earth isotopes (see Experimental Section). With a resolution (fwhm) of about 2.3 keV for the Ge(Li) detector used in the experiments, simultaneous counting of two or more of the listed rare-earth isotopes was easily achieved. The relative extraction efficiencies of the lanthanides were found to depend on the shaking time as well as the composition of the organic phase. For example, in the case of the La3+-Lu3+competition at pH 6, the relative extraction efficiency of Lu3+/La3+reached a maximum value of 6.0 after 2 min of shaking with an organic phase consisting of 20% by volume of heptanol in chloroform. Thm ratio tends to decrease by 10% when shaking time was 10 min. The composition of the organic phase appears to be an even more important factor in determining the relative extraction efficiency. Figure 3 shows the relative extraction efficiency of Lu3+/La3+as a function of solvent composition for mixtures of chloroform and heptanol. In pure chloroform, the relative extraction efficiency is a factor of 3.7 favoring Lu3+over La3+ but reaches a maximum value of 6.0when the heptanol concentration increases to 20% by volume. In pure heptanol, no preference for the extraction of Lu3+over La3+was observed. The effect of solvent on the relative extraction efficiencies of the lanthanides has been observed in neutral crown ether systems (I). The relative extraction efficiencies of other lanthanides including P P ,Sm3+,Eu3+,Tb3+,Er3+,and Yb3+ compared to Lu3+ in a system with 20% heptanol in the organic phase is shown in Figure 4. The extraction selectivity appears to change more rapidly from La3+to Eu3+but seems to level off from Dy3+ to Lu3+. The ionic radii of the lanthanides are known to vary from 1.15 A for La3+to 0.95 A for Lu3+. The cavity diameter of sym-dibenzo-16-crown-5-oxyacetic acid has been estimated to be about 2.0-2.2 A (3). It
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Z Figure 4. Extraction efficiencies of seven lanthanides relative to Lu3+ with sym-dibenzo-16-crown-5-oxyacetic acid in a mixture of 80:20 chloroform-heptanol. Lanthanides l i s t e d i n increasing
appears that all lanthanide cations fit the cavity of the crown ether carboxylate well, and the selectivity may be controlled by factors other than cavity size alone. Radioisotope tracer experiments performed with synthetic seawater and natural river water samples also indicated good extraction efficiencies (>95%) for the spiked lanthanide ions into an organic phase of 8020 chloroform-heptanol that acid. The high contained sym-dibenzo-16-crown-5-oxyacetic extraction efficiencies for the lanthanides observed in such complex aqueous systems suggest potential applications of the crown ether carboxylic acid in analytical chemistry as well as in metallurgy, which involves solvent extraction or membrane transport of the rare-earth elements. Registry No. La, 7439-91-0; Pr, 7440-10-0; Sm,7440-19-9; Eu, 7440-53-1;Dy, 7429-91-6;Er, 7440-52-0;Yb, 7440-64-4;Lu, 7439-94-3; syn-dibenzo-16-crown-5-oxyacetic acid, 78708-41-5. LITERATURE CITED (1) Masake, Y.; Hdeyuki, N. Anal. Lett. 1982, 75, 1197. (2) Kolthoff, I. M. Anal. Chem. 1979, 51, 1R. (3) Strzelbicki. J.; Bartsch, R. A. Anal. Chem. 1981, 53, 1894. (4) Strzelbicki, J.; Bartsch, R. A. Anal. Chem. 1981, 53. 2247. (5) Kang, S. I.; Czech, A,; Czech, 6. P.; Stewart, L. E.; Bartsch, R. A. Anal. Chem. 1985, 57, 1713. (6) Bartsch, R. A.; Heo, 0.S.; Kang, S. I.; Liu, Y.; Strzelbickl, J. J . Org. Chem. 1982, 4 7 , 457. (7) Mok, W. M.; Shah, N. K.; Wai, C. M. Anal. Chem. 1988, 58, 110. (8) Bunzli, J. C. G.; Oanh, H. T.; Gillet, B. Inorg. Chim. Acta 1981, 53,
L219. (9) Simon, J. D.; Moomaw, W. R.; Ceckler, T. M. J . Phys. Chem. 1985, 89, 5659. (10) Bunzli, J. C. G.; Wessner, D. Isr. J . Chem. 1984, 2 4 , 313.
Jian T a n g M.Wai*
C. Department of Chemistry University of Idaho Moscow, Idaho 83843
RECEIVED for review May 6,1986.Accepted August 11,1986.
Oxidative Mechanism of Ascorbic Acid at Glassy Carbon Electrodes Sir: In our previous reports (I+?), it was shown that a glassy carbon electrode (GCE) after vacuum heat treatment (VHT) (1)or properly polishing under strictly clean conditions (2, 3) is highly active for the electrooxidation of ascorbic acid.
These treatments not only reduce the overpotentialof ascorbic acid oxidation at a W E by more than 300 mV but also double the magnitude of the cyclic voltammetric peak current for this totally irreversible electrooxidation. This increase in the peak
0003-2700/88/035&3235$01.50/0 0 1986 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 58, NO. 14, DECEMBER 1986
current indicates that there is a change in the mechanism of this oxidation when the electrode is activated. An understanding of the electron transfer mechanism at an inactive and an active GCE is important for appreciation of the activation mechanism with a GCE. Results from our previous work (1-3) indicated that the electron transfer of the electrooxidation of ascorbic acid at a GCE may possibly occur through three types of sites on a GCE surface. One of these sites is the “pristine” carbon, which can be created by VHT (1). The second is the electroactive quinoidal-like surface functionalities (4),which can be produced by means of electrochemical oxidation (5, 6 ) , radio frequency plasma treatment ( 7 , 8 ) , and chemical oxidation (9). The third type is the sites that are blocked by impurities (2,3)or electroinactive surface oxide functionalities (10). The activities of these three sites for the electron transfer differ with respect to one another and depend on the electrode potential. The pristine carbon appears to be the most probable path for the electron transfer for the oxidation of ascorbic acid ( 1 ) a t lower potentials (Le., 200-250 mV vs. Ag/AgCl (saturated KC1) at pH 2.0). As the electrode potential approaches the redox potential (ca. 250 mV a t pH 2.0) of the electroactive surface functionalities, these functionalities may mediate the charge transfer (5, 7 , 8 ) . Finally, at sufficiently high potentials, the blocked sites begin to participate in the electron transfer. Although the results of our previous studies have led us to propose the above model of the electron transfer at GCE’s, details of the electron transfer mechanism have not been provided. In this paper, the oxidative mechanism of ascorbic acid at a GCE is studied and is used to verify the above-proposed electron transfer model. In this study, the techniques of cyclic voltammetry and semiintegral (or convolution) cyclic voltammetry are used to evaluate the kinetic parameters such as transfer coefficients and number of transferred electrons. The electrooxidation mechanisms of ascorbic acid a t an activated GCE with a “pristine” carbon surface, a t a partially activated GCE, and at an inactive GCE are studied.
EXPERIMENTAL SECTION Reagent grade ascorbic acid (J. T. Baker Chemical Co., Philipsburg, NJ), KCI, NaOH, and H3P04(Mallinckrodt, Inc., Paris, KY) were used without further purification. All solutions were prepared daily with NANOpure water (Sybron Barnstead, Boston, MA). Glassware was cleaned in an 2-propanol/KOH bath followed by immersion in 1M H2S04and was thoroughly rinsed in NANOpure water. All solutions for electrochemical experiments were purged with prepurified nitrogen. The glassy carbon electrodes were cut from a plate of Tokai GC-20 (Tokai Carbon Co., Ltd., Tokyo, Japan). The geometric area of each GCE was 0.090 f 0.012 cmz. The electrode was polished according t o the procedures described previously (2). The vacuum heat treatment method used to create a “pristine” carbon surface was the same as that described in a previous paper (1).
The electrochemical measurements were conducted with an in-house-builtthree-electrode potentiostat. The potentiostat was interfaced to an Apple IIe computer using a 12-bit D/A and A/D interface card (Interactive Microwave Inc., State College, PA). The resolution of synthesized excitation signals from this interface was 0.48 mV. Software for cyclic voltammetry, chronocoulometry, and semiintegral cyclic voltammetry was written in machine language t o provide optimal speed for instrumentation control, data acquisition, real-time data display, and data analysis. The “G-I” algorithm (11) was used in the semiintegration of cyclic voltammetric data.
RESULTS AND DISCUSSION The electrooxidative mechanism of ascorbic acid at a mercury electrode has been extensively studied by Ono et al. (12),Perone et al. (13),and Ruiz et al. (14,15). Although the mechanism of this reaction at above pH 8 is still disputed (16),
a mechanism proposed by Ruiz (14, 15) for the reaction at below pH 8 is widely accepted (16). Figure 1 illustrates the scheme of the mechanism. This mechanism involves a predissociation of a proton to give the monoanionic species followed by a 1 e-, 1 H+ oxidation of the monoanionic species to form a radical anion, which then undergoes a second irreversible 1 e- oxidation to dehydroascorbic acid. The latter species is rapidly protonated and then dehydrated to yield the final product of 2,3-diketogulonic acid. The hydration rate at pH 7.2 was estimated by Perone (13)to be about 1.2 x lo3 s-l. On the basis of the above mechanism, if one assumes that the proton transfer step is not the rate-determining step (rds) and the overall reaction is totally irreversible, the rate constant for the overall reaction can be expressed (17) as kobsd = kZk3/(k-Z + k3) (1) where k1 and k+ are the forward and backward rate constants of the first electron transfer step, respectively, and k3 is the standard rate constant for the second electron transfer step. By replacing all of the rate constants with potential-dependent rate constants and assuming a transfer coefficient of 0 = 0.5 for each electron transfer step, eq 1 becomes ( 17) l/k(EappJ = l / W , exp(0.5FAEl/RT)1 + ~ / W * ~Z X P ( ~ . ~ F G Z / R(2) T)I k*2 = k o 2 e ~ p [ 0 . 5 F ( E ” -’ ~ED’,)/RT] ~ (3)
where k(EnPpJ is the apparent rate consatnt for the overall reaction at an applied potential of Eappl; kD1and k ” , are the apparent rate constants for the first and second steps, re- E”,; AElz is .Eapplspectively; AE,is Eappl and is the formal potential for the overall reaction. As indicated from eq 1 and 2, if the first electron transfer is the rds (i.e,, k3 >> k-2) k(EappI)= k o l exp(0.5FAEl/R‘I?
(4)
or if the second electron transfer is the rds (i.e., k 3
