Kinetic studies of dioxygen and superoxide ion in acetonitrile at gold electrodes using ultrafast cyclic voltammetry

Journal of

Electroanalytical Chemistry Journal of Electroanalytical Chemistry 576 (2005) 95–103 www.elsevier.com/locate/jelechem

Kinetic studies of dioxygen and superoxide ion in acetonitrile at gold electrodes using ultrafast cyclic voltammetry Zhiyong Guo, Xiangqin Lin

*

Department of Chemistry, University of Science and Technology of China, Hefei, Anhui 230026, PR China Received 26 April 2004; received in revised form 11 October 2004; accepted 12 October 2004 Available online 24 November 2004

Abstract On the basis of a circuit with ohmic drop compensation by electronic positive feedback, ultrafast cyclic voltammetry was used to study the electrochemical behavior of dioxygen and superoxide ion at gold electrodes in acetonitrile containing 0.9 M tetraethylam monium tetrafluoroborate as the supporting electrolyte. At low scan rates, the O2 =O
2 couple is controlled by diffusion, and its redox process is an irreversible one with the rate constant (7.38 ± 1.25) · 10
3 cm s
1. The diffusion coefficient of O2 in such a solution was deduced as (4.87 ± 0.08) · 10
5 cm2 s
1, while at high scan rates, it is controlled by the adsorbed O2 at a rate constant  2
(5.21 ± 0.66) · 104 s
1. At moderate scan rates, it is controlled by both of them simultaneously. As to the O
2 =O2 couple, its faradaic signal was greatly distorted at low scan rates. However, when the scan rate was increased sufficiently, the perturbing reactions would be Ôout-runÕ by a transient experiment. Then, kinetic information could be obtained from the well-shaped voltammograms. In  4
1 our experiments, the adsorption of O
2 was found, and the rate constant was deduced as (1.95 ± 0.28) · 10 s .
2004 Elsevier B.V. All rights reserved. Keywords: Ultrafast cyclic voltammetry; Dioxygen; Superoxide; Peroxide; Ultramicroelectrodes

1. Introduction The electrochemical study of dioxygen has a long history [1]. It is still of great interest now in this type of research, because dioxygen is vital to life and nature, and  the electrogenerated superoxide ion O
2 , which could subsequently generate other active oxygen species, relates to the aging process and many maladies according to research into free radicals [2]. The electrochemical investigation of dioxygen has been carried out in a wide variety of media. A first  is medium is alkaline aqueous solution. Because O
2 a strong nucleophile and involved in a disproportionation process with fast second order rate constants between 107 and 1010 M
1 s
1 in water [3], appropriate

*

Tel.: +86 551 3606646; fax: +86 551 3601592. E-mail address: [email protected] (X. Lin).

0022-0728/$ - see front matter
2004 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2004.10.011

surfactants should be present such as isoquinoline [4,5], triphenylphosphine oxide [6,7], a-quinoline [8,9] and sodium dodecyl sulfate [10], etc. A second is the range of aprotic media usually used [11–28, and references therein], such as N,N-dimethylformamide (DMF), dimethyl sulfoxide (DMSO) and acetonitrile (AN), in which DMSO is the most widely used. Compared with water, they have wider negative potential windows so as to avoid solvent breakdown as a competing reaction, and a significantly greater solubility of dioxygen. In the study of the dioxygen–superoxide system in non-aqueous aprotic solvents, various electrode materials have been applied such as carbon, platinum, gold, mercury, noble metal–mercury alloys, etc [11–28]. In general, there is agreement that the reduction of dioxygen in aprotic media occurs usually in two principal steps, reversible/quasi-reversible/irreversible and irreversible, respectively. The first is a one-electron reduction

96

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

of dioxygen to superoxide, and the second is a successive one-electron process forming the peroxide anion, which may react with the solvent [17,18]. The formal potentials of the redox couples are related to the solvents and electrode materials used [21]. O þ e
! O
 ð1Þ 2

2

 2

O
2 þ e ! O2

ð2Þ

The first step is conventionally regarded as a simply diffusion-limited one. However, the adsorption of dioxygen on the electrode surface has also been discussed [19,20]. Compton and co-workers recently reported [24] that, in the electrochemical reduction process of dissolved dioxygen in DMSO with tetraethylammonium perchlorate (TEAP) as the supporting electrolyte on Au microelectrodes, an adsorption intermediate of dioxygen was found in transient potential step chronoamperometric studies, with the rate constant k1480 s
1 for the electroreduction of dioxygen adsorbed on the gold electrode in 0.2 M TEAP + DMSO. A monolayer of dioxygen was composed of 2 · 1014 molecules cm
2. K

k1

 O2 ! O2 ðadsÞ ! O
2 slow

ð3Þ

There have been only a few studies [11,16,18, 19,23,24,27] of the second step, but kinetic parameters have not been reported to our knowledge, because it is complicated by several factors. The first is the dismuta tion process of O
2 itself.  2
2O
2 ! O2 þ O2

ð4Þ

The second factor is that the peroxide formed ion may react with the solvent. The third and the most important factor is that it is difficult to achieve an absolutely aprotic solvent experimentally because moisture (a proton source) is one of the most difficult impurities to eliminate from non-aqueous solvents, and water concentrations of the order of 5 · 10
4 M (10 ppm), only one order of magnitude lower than the dioxygen concentration, are often present in the best purified solvents. Thus, some following reactions of the electrogenerated  2
O
2 and O2 would occur due to the presence of protons [11–28]:

All reactions in the above sequence take place so rapidly on the time scale of low scan rates that no anodic peak corresponding to the oxidation of peroxide ion to superoxide could be found in the reverse scan. As a useful tool in the pursuit of kinetics, cyclic voltammetry could provide a large amount of electrochemical information such as: kinetic processes for heterogeneous electron transfer reactions, coupled homogeneous chemical reactions, electrogenerated transient intermediates, etc [29]. Obviously, the more rapid scan rate t employed, the smaller time window of the experiments would be. If t is high enough, the above-mentioned sequence of reactions will be Ôout 2
couple could be easily runÕ, i.e., the O
2 =O2 examined.  Studying the electrochemical behavior of O2 and O
2 in non-aqueous media is actually fundamental work for determination in biological systems. In this approach, ultramicroelectrodes showed special advantages [24,27,28,30–34]. The basic advantage is probably that very limited physical and chemical perturbations could be effected on the biological system being measured. The use of ultramicroelectrodes allows the application of fast potential scanning, which could provide additional advantages such as improving the stability of the working electrode [35], decreasing the diffusion layer thickness to avoid being blocked by tissues in in-vivo experiments [36] and gratifying the need for rapid response. However, most of the measurements were conducted at conventional scan rates for technical simplicity. In this work, to demonstrate the above hypothesis and study the adsorption behavior of the first step better, ultrafast cyclic voltammetry was employed by virtue of a homemade circuit with online ohmic drop compensation by electronic positive feedback [37]. Due to the very high scan rate applied, special experimental conditions, i.e., a very high concentration of the supporting electrolyte (0.9 M) was used as suggested by Amatore et al. [38] to avoid the coupling of the diffusion layer and the double layer.

2. Experimental

 þ O
2 þ H ! HO2

ð5Þ

2.1. Reagents

2HO2 ! O2 þ H2 O2

ð6Þ

HO2 þ e
! HO
2

ð7Þ


þ O2
2 þ H ! HO2

ð8Þ


þ
HO
2 þ H þ 2e ! 2OH

ð9Þ

Chemical reagents used in the experiments were all obtained in the highest grades commercially available. Tetraethylammonium tetrafluoroborate (NEt4BF4) purchased from Aldrich was dried at 60 C under vacuum, and stored in a vacuum desiccator. Acetonitrile (HPLC quality, Shandong Yuwang Industrial & Commercial Corporation, China) was twice distilled on calcium hydride just before use. The purity of nitrogen and dioxygen (Shanghai Pujiang Special Gas Corporation, China)



HO
2 þ H2 O þ 2e ! 3OH

ð10Þ

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

was more than 99.999%. Anthracene (Sigma) was recrystallized in acetonitrile before use. 2.2. Apparatus For electrochemical experiments at low scan rates, a CHI600A voltammetric analyzer (CH Instruments, USA) was employed. For ultrafast cyclic voltammetric experiments, a home-built system was utilized with an online positive feedback IR compensation. The input signals were generated from a Model G305 waveform generator (10 MHz, HUNG CHANG, Korea), which then passed through the electrochemical cell and a homemade circuit. Finally, the voltammograms were recorded with a Tektronix Model TDS220 digitizing oscilloscope with 100 MHz band-pass. Data were downloaded and stored on a Pentium IV computer for further analysis. A 10 MHz @ 0.1 dB bandwidth of this circuit gave good results in transient electrochemical studies. Tests have shown that this circuit could record undistorted voltammograms up to 2.67 MV s
1 scan rate, and 3 MV s
1 could be reached if an error of 30 mV was tolerated. Its integrated theoretical behavior and experimental demonstration of its validity have been reported previously [37]. A two-electrode system was employed, including a 5 lm gold microdisk electrode as the working electrode, a platinum wire counter/quasi-reference electrode and a 200 ll glass cell mounted directly on the printed circuit board. The Au microdisk working electrode comprised a Au wire (Goodfellow Metals, 5 lm nominal diameter, purity 99.99%) sealed in glass and was carefully polished using successively 0.3 and 0.05 lm alumina powder (CH Instruments, USA) on a clean, damp polishing cloth (CH Instruments, USA). It was then sonicated continuously in distilled deionized water and acetonitrile each time just prior to recording voltammograms. Its diameter was calibrated electrochemically using 5 mM K3Fe(CN)4 in 0.5 M KCl with a diffusion coefficient of 6.3 · 10
6 cm2 s
1 [39], and the value was 4.0 ± 0.2 lm. Another gold electrode (4 mm diameter, Tianjin Lanlike, China) was used to perform experiments at low scan rates. A platinum wire (1 mm diameter) was used as the counter/quasi-reference electrode, whose surface area was about 0.3 cm2 and 106 times larger than that of the Au microdisk working electrode. As it was postulated and observed [40] that its use might lead to interferences by eventual dissolution and subsequent adsorption of Pt on the Au working electrode, the solution was frequently replaced to eliminate this problem. The electrochemical redox reaction of anthracene was used to calibrate the potential axis as an internal reference because of its high reversibility suitable for various scan rates, at which the conventional SCE might not be used. However, the anthracene could not be added in the solution for potential calibration of the oxygen

97

waves since the reduction of anthracene might obscure one of the oxygen waves of this system. In this case, the calibration was conducted for the bulk solution. Our result showed that the potential floating due to the Pt quasi-reference was small in the ultrafast CV experiments. Thus, this calibration should be valid and should not interfere significantly with the extraction of kinetic parameters. 2.3. Procedures Cyclic voltammetry was carried out starting with the negative scan and then reversal of the scan unless otherwise stated. The amount of residual water in the acetonitrile was estimated roughly as 0.4–0.5 mM from the slight bump of the background voltammogram in 0.9 M NEt4BF4 + AN solution containing no O2. Due to the high scan rate employed, the time lag tlag caused by circuit distortion and ÔmechanicalÕ delay should be considered. Fortunately, it did not alter the voltammetric shapes, so it could easily be corrected by translating the potential axis +ttlag in the negative scans and
ttlag in the positive scans [38]. In the determination of kinetic parameters, a 10 ns tlag was corrected. The electrochemical solution was bubbled for 10 min with AN saturated O2 gas to ensure O2 saturation before measurements. A flow of O2 was maintained over the solution during experiments. However, AN saturated nitrogen gas was used to remove O2 when obtaining background voltammograms. All experiments were conducted at room temperature (20 ± 1 C).

3. Results and discussion  3.1. Electrochemical behavior of the O2 =O
2 couple in acetonitrile at gold electrodes A series of voltammograms recorded in O2 saturated 0.9 M NEt4BF4 + AN at the 5 lm working electrode at different scan rates is shown in Fig. 1(A). Their shapes were similar and independent of the scan rates, suggesting that typical steady-state electrochemical behavior occurred up to 0.1 V s
1. A well-defined limiting current plateau IL can be clearly seen. For a microdisk electrode, IL can be expressed as in the generally accepted equation: I L ¼ 4nFDc0 r

ð11Þ

where all symbols have their usual significance. Introducing n = 1, c0 = 8.2 mM [22], r = 2.0 lm and IL of each voltammogram into the equation, the diffusion coefficient of dioxygen DO2 in 0.9 M NEt4BF4 + AN was calculated as (4.87 ± 0.08) · 10
5 cm2 s
1, as shown in the inset of Fig. 1(A). This value is a little smaller than that reported as 7.12 · 10
5 cm2 s
1 in 0.1 M NEt4ClO4

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

0 (A)

(B) cm 2 s

-20

5

E / V vs. Pt

-10

-0.7 0.190 0.007

-0.8 -0.9

D/

I / nA

1

4.87 0.08

dE/dlg[ (IL I)/I]

98

-30

-1.0

-1.0

-0.5

0.0

-0.5 0.0

1.0

0 (D)

-0.1 (C) -0.2 0.00738 0.00125

-0.3 -0.4

0.306 0.011

-0.5 ks / cms

-2

0.5

lg[(IL I)/I]

0

I / nA

0

E E / V vs. Pt

E / V vs. Pt

-10 -20 -30

2

-1

ln[(IL I)/I (IL I )/I ]

0

E / V vs. Pt

Fig. 1. The electroreduction behavior of saturated O2 in AN + 0.9 M NEt4BF4 solution on a 5 lm gold electrode. (A) A set of steady-state voltammograms at different scan rates: 0.03 V s
1 (solid rectangles), 0.05 V s
1 (solid line), 0.06 V s
1 (open triangles), 0.08 V s
1 (open cycles), 0.1 V s
1 (open rectangles); inset: the error bar of DO2 determined. (B) The plot of E vs. log [(IL
I)/I] of the set of voltammograms, the symbols 0 are same as the above; inset: the error bar of the slope. (C) The plot of E
E0 vs. ln[(IL
I)/I
(IL
Is)/Is] of the set of voltammograms, the symbols are same as the above; inset: the error bars of a and ks. (D) The voltammogram of 0.05 V s
1 obtained experimentally (solid line) and the simulated curve (open circles) with the following parameters: a = 0.306, ks = 7.38 · 10
3cm s
1, DO2 = 4.87 · 10
5 cm2 s
1, r = 2 lm, E0 =
0.55 V vs. the Pt reference.

(TEAP) [22] and 7.05 · 10
5 cm2 s
1 in 0.1 M NBu4ClO4 (TBAP) [23]. The difference in the DO2 values obtained may be generated from the different viscosities in different concentrations of different types of supporting electrolyte. The viscosity in 1 M solution would be 10–20% higher than in 0.1 M solution, and a value of 2%/C temperature coefficient should also be taken into account for the diffusion coefficient obtained [41]. According to WaldenÕs Rule, the diffusion coefficient is the reciprocal of the viscosity. After these consideration, the DO2 value we obtained is in good agreement with the above literature values. However, a value of 11 · 10
5 cm2 s
1 was reported for the 0.1 M TEAP/AN system [33], which is much larger than our value, even after consideration of all above-mentioned factors. Sawyer and Roberts [16] reported that the diffusion  coefficient of O
2 is significantly influenced by the ion pairing with the supporting electrolyte cations, and DO2  3:5  DO
2  was found in the TEAP/DMSO system. Assuming that the solvation, if any, of the superoxide ion should not be radically different in AN and in DMSO, and that the cations are same, thus DO
2  could be roughly estimated as 1.39 · 10
5 cm2 s
1. Fig. 1(B) shows the plot of E vs. log[(IL
I)/I] of the series of the voltammograms shown in Fig. 1(A). The slope dE/dlog[(IL
I)/I] as indicated in the inset of Fig.

1(B) is 0.190 ± 0.007 which is absolutely consistent with the theoretical slope of an irreversible wave 2.303RT/ anF = 0.190 (a = 0.306, vide infra), suggesting that the  couple in AN electrochemical process of the O2 =O
2 on gold electrodes is an irreversible one. The fact that the intercepts of these lines varied in a small way indicated the potential floating caused by using a quasireference electrode. Fortunately, such a floating was not serious and did not interfere with the extraction of kinetic parameters greatly. As reported by Galus et al. [42], steady-state voltammogram could be used to determine the heterogeneous electron transfer rate constant ks. When the oxidized form only is present initially in the solution, the voltammetric response of a quasi-reversible reduction reaction under steady-state conditions on a microdisk electrode is given as follows (see Appendix A): 0

ð4DO =pk s rÞ exp½anf ðE
E0 Þ ¼ ðI L
IÞ=I
ðI L
I s Þ=I s

ð12Þ

and 0

E ¼ E0 þ ð1=nf Þ lnðDR =DO Þ þ ð1=nf Þ ln½ðI L
I s Þ=I s  ð13Þ 00

where E is formal potential of the redox couple, E is the potential of the kinetically controlled process, I is

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

the current of the kinetically controlled process measured at potential E, Is is the reversible current calculated with Eq. (13), n is the electron number, f = F/RT, D is the diffusion coefficient, ks is the heterogeneous electron transfer rate constant, r is the electrode radius and other symbols have their usual significance. Thus, the 0 plot of E
E0 vs. ln[(IL
I)/I
(IL
Is)/Is] will be linear with the slope equal to 1/anf and the intercept (
1/anf) ln (4DO/pksr), as shown in Fig. 1(C). Then ks could be easily calculated by dividing the intercept by the slope, if r and DO are known. In this case, ks = (7.38 ± 1.25) · 10
3 cm s
1 and a = 0.306 ± 0.011 according to the slope, as shown in the inset of Fig. 1(C). As suggested by Wu and Zhang [43], 4DO/ pksr > 30 is a criterion of an irreversible steady-state reduction wave. In our case this value was equal to 42, which suggested that the electrode reaction is an irreversible process again. Sawyer et al. [21] have reported that the peak-to-peak potential separation DEp of the reduction of dioxygen to superoxide ion on a Au electrode in 0.1 M TEAP + AN solution at a scan rate of 0.1 V s
1 was 470 mV, thus ks could be deduced as about 1.1 · 10
4 cm s
1 according to NicholsonÕs method [44], which is much lower than our value. Indeed, ks is a function of various factors such as the temperature, the solvent used, the supporting electrolyte used and its

99

concentration, electrode material and the electrode surface pre-treatment, etc. Actually, the study by Collins et al. [45] implied that the Au electrode used by Sawyer et al. was perhaps not a recently polished one, but a one that had experienced some cyclic sweeps, which would greatly increase DEp and thus greatly decrease the calculated ks accordingly. To test the reliability of these deduced parameters, a voltammogram was simulated [46] with the following parameters: a = 0.306, ks = 7.38 · 10
3 cm s
1,
5 2
1 00 DO2 ¼ 4:87  10 cm s , r = 2 lm, E =
0.55 V vs. the Pt reference. As shown in Fig. 1(D), there was a perfect fit between the experimental and the simulated curves, which demonstrated clearly the validity of the results. It is difficult to obtain kinetic information from a voltammogram including two components, diffusion and adsorption. Usually at low scan rates, the peak current is controlled by diffusion, and at high scan rates, by adsorption, as the peak currents are proportional to t1/2 and t, respectively. To study the adsorption of dioxygen at Au electrodes, ultrafast cyclic voltammetry was used. The results are shown in Fig. 2. Fig. 2(A) shows the voltammogram of the reduction of O2 at the 5 lm working electrode in AN + 0.9 M NEt4BF4 solution at a scan rate of 30 kV s
1. It is

Fig. 2. Electrosorption behavior of oxygen in O2 saturated 0.9 M NEt4BF4 + AN solution on a 5 lm gold electrode. (A) Voltammogram at 30 kV s
1 (solid line) and voltammograms in blank solution at the transferred electrode without cleaning (open circles) and cleaned (dotted line) at 30 kV s
1, after a cycle in O2 saturated 0.9 M NEt4BF4 + AN solution. (B) Voltammograms of O2 saturated 0.9 M NEt4BF4 + AN solution at 12 kV s
1 (open circles), 30 kV s
1 (open rectangles), 48 kV s
1 (solid line), 75 kV s
1 (open triangles) and 120 kV s
1 (solid rectangles) and the simulated curve (dotted line) for a diffusion controlled redox reaction in which Butler–Volmer kinetics were obeyed and following parameters were used: c0 = 8.2 mM, r = 2.0 lm, DO2 ¼ 4:87  10
5 cm2 s
1 ; DO
2  ¼ 1:39  10
5 cm2 s
1 , ks = 7.38 · 10
3 cm s
1, a = 0.306, n = 1, Ru = 0 X, Cdl = 6 pF, t = 48 kV s
1, 0 E0 =
0.55 V vs. Pt; inset: comparison between the experimental (solid line) and the simulated (dotted line) results for a scan rate of 48 kV s
1. (C) The plot of Ipc vs. t; (D) plots of DEp (potential separation between anodic and cathodic peaks) vs. t: the experimental data with 10 ns time lag correction (solid rectangles), the theoretical prediction for kad = 5.21 · 104 s
1 and a = 0.5 (solid curve) and error bar of kad (the inset).

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

composed of bell-shaped peaks, which is characteristic of an electron transfer for an adsorbed layer of redox centers. No faradaic peaks appeared in the background voltammogram at the same scan rate with a newly polished electrode, but with an electrode which was not polished after it had experienced a cycle in a saturated O2 0.9 M NEt4BF4 + AN solution, faradaic peaks were present. This suggests the existence of adsorption behavior. Fig. 2(B) shows a set of voltammograms under the same experimental conditions as above at different scan rates. The peaks became a little slanted when t was increased, suggesting that the voltammetric time scale s  RT/Ft became smaller, viz., ks became comparable with 1. To demonstrate further the participation of an adsorbed species in the heterogeneous electron-transfer process, a diffusion-controlled voltammogram was simulated [47] under the same conditions as the above experiment using the parameters as follows: c0 = 8.2 mM, r = 2.0 lm, DO2 ¼ 4:87  10
5 cm2 s
1 ; DO
2  ¼ 1:39  10
5 cm2 s
1 , ks = 7.38 · 10
3 cm s
1, a = 0.306, 0 n = 1, Ru = 0 X, Cdl = 6 pF, t = 48 kV s
1, E0 =
0.55 V vs. the Pt reference. From the inset of Fig. 2(B), by comparing the simulated and the experimental voltammograms at a scan rate of 48 kV s
1, we found that the peak current of the simulated was much lower than that of the experimental voltammogram and DEp was much larger, which clearly indicated that the process at high scan rate was controlled not by diffusion but by adsorption. Fig. 2(C) is the plot of Ipc vs. t. It exhibits an apparent linear trend with the line through the origin, which indicates that the process was adsorption-controlled again. According to the equation [29] I p ¼ n2 F 2 tAC =4RT

ð14Þ

where C* can be obtained as (1.82 ± 0.07) · 10
6 mol m
2, i.e., (1.1 ± 0.04) · 1014 molecules cm
2 assum˚ . This ing that the diameter of dioxygen is about 10 A value is nearly consistent with that reported [24] (2 ± 1.5) · 1014 molecules cm
2 in 0.2 M TEAP + DMSO. The difference might be due to the fact that the Au electrode was a Au (1 1 1) single crystal in the experiments performed by Compton and co-workers [24], while it was a polycrystalline one in our case. The quantity of charge, Q, required to reduce adsorbed O2, was 17.6 lC cm
2, which was larger than the literature value (Q < 15 lC cm
2) [48]. This was understandable because the solvent used in the literature was 0.05 M sulfuric acid solution, while ours was acetonitrile, in which the adsorption of O2 was very much more marked [20]. In the above-mentioned experiments, an increase of DEp with the scan rate, caused by the peaks shifting toward higher potentials, was observed. This was due to the increasingly smaller time scale of the voltammetric experiment vis-a`-vis the electron transfer rate constant,

which made ks  1 provided that the scan rate was sufficiently high. The increase of DEp is a characteristic of the transition to a kinetic regime where the electron transfer becomes limiting and governs the voltammetric behavior. Such a transition between fast and slow electron transfer regimes allows therefore a facile and precise determination of k. Assuming that for a oneelectron transfer, DEp is greater than 200 mV, k could be calculated with the help of the following equation [49], log k ¼ a logð1
aÞ þ ð1
aÞ log a
logðRT =nF tÞ
að1
aÞnF DEp =2:3RT :

ð15Þ

Assuming a = 0.5, we obtained the rate constant of the reduction of the adsorbed dioxygen kad = (5.21 ± 0.66) · 104 s
1 from the data in Fig. 2(D), which revealed the fast charge transfer rate of the process. This value was about 100 times higher than the reported value of 500 s
1 [24] in 0.2 M TEAP + DMSO, mainly due to the fact that the solvent dependence of dioxygen adsorption is AN  DMSO > DMF [20]. Obviously at moderate scan rates, the redox process  couple was controlled simultaneously of the O2 =O
2 by diffusion and adsorption. Thus, it could be predicted that the faradaic signals corresponding to the respective modes could be found simultaneously in the voltammogram at moderate scan rate if only the adsorption were strong enough. This has been demonstrated clearly by our experimental results. In the voltammogram at a scan rate of 4.8 kV s
1 shown in Fig. 3, two pairs of peaks appeared. A diffusion-controlled voltammogram was simulated and is also shown in Fig. 3, with the same parameters as those of Fig. 2(B) except for t = 4.8 kV s
1. By comparing the experimental with the simulated voltammogram, it is obvious that peaks b and d correspond to the diffusion behavior, while peaks a and c correspond to adsorption. It is generally assumed that the post-peaks for the adsorbed dioxygen should

c

d

0.0

I/ A

100

b a -0.1 -1

E / V vs. Pt Fig. 3. Experimental voltammogram (solid line) of O2 saturated 0.9 M NEt4BF4 + AN solution at a 5 lm gold electrode at 4.8 kV s
1 and a simulated curve (dotted line) with parameters as in Fig. 2(B) except for t.

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

both be on the negative side of the dioxygen diffusioncontrolled peaks [29], but in our case they are on the other side. Apparently, this is due to the effect of the kinetics, i.e., the diffusion-controlled reduction of O2  with a slow electron transfer rate while the to O
2 adsorption-controlled kinetics are fast.  2
3.2. Electrochemical behavior of the O
2 =O2 couple in acetonitrile at gold electrodes  2
The electrochemical behavior of the O
2 =O2 couple can be studied if we extend the potential window to a more negative position. However, due to the abovementioned perturbing chemical and electrochemical reactions (Eqs. (5)–(10)), the cathodic peak correspond 2
ing to the reduction of O
2 to O2 was greatly distorted and no anodic peak was found corresponding to the
 reoxidation of O2
2 to O2 at low scan rates on a conventional electrode, as is shown in the inset of Fig. 4(A). However, from the voltammogram at 48 kV s
1 in Fig. 4(A), we found that peaks b and c corresponding  to O2
were in good shape, to the reduction of O
2 2 which suggested that those reactions leading to distortion were successfully Ôout-runÕ by increasing the scan rate. No faradaic peaks appeared in the background voltammogram with the electrode newly polished, but they did appear with the electrode unpolished after it had experienced a cycle in a saturated O2 0.9 M NEt4B-

I/ A

I/ A

0.3 (B)

c 0 -1

F4 + AN solution. Thus, this demonstrated the existence of adsorption. To show the evolution of the second reduction wave and determine the low limit of the scan rate at which the above-mentioned reactions could be Ôout-runÕ, a series of voltammograms from low to medium scan rates was recorded, as shown in Fig. 4(B). At 3 kV s
1, no anodic peak could be seen at all, while at 12 kV s
1, an anodic peak could be seen with the shape distorted to a certain extent. When the scan rate was increased up to 18 kV s
1, a better-shaped anodic peak could be seen. The lifetime of the electrochemically generated O2
is 2 estimated as about 10
4 s under the experimental conditions. The peaks were bell-shaped, characteristic of adsorption. A set of voltammograms at scan rates higher than 18 kV s
1 are shown in Fig. 4(C). A linear relation between Ipc and t exists as is illustrated in the inset of  Fig. 4(C), clearly demonstrating that O
2 was adsorbed on the surface of the electrode also. By the above could be calculated as mentioned method, C* of O
2
7
2 (3.06 ± 0.18) · 10 mol m , much lower than that of dioxygen, perhaps due to the negative electrode charge. Using the data in Fig. 4(D), we could extract the rate  2
couple as (1.95 ± 0.28) · 104 constant of the O
2 =O2
1 s . A previous paper reported that no adsorption of  O
2 was apparent [20], we think this was because the

d

1 (A)

b

0.0

0.1 mA

a

-2

-2

-1

E / V vs. Pt

0

-0.3 -2.5

0

-2.0

-1.5

E / V vs. Pt

E / V vs. Pt 1 (C)

0

Ep / V

(D)

0.4

Ipc / A

Ipc / A

101

1.95

0.28

kad /

10 s

0.3

0.2

0.0

-1

0.4

40.0k

80.0k

/Vs

-2

E / V vs. Pt

-1

20k

4

1

40k 60k

/Vs

1

 2
Fig. 4. Electrochemical behavior of the O
2 =O2 couple on a 5 lm gold electrode in 0.9 M NEt4BF4 + AN solution. (A) Voltammogram on a 5 lm gold electrode at 48 kV s
1, and voltammograms in blank solution at the transferred electrode without cleaning (open circles) and cleaned (dotted line) at 48 kV s
1, after a cycle in O2 saturated 0.9 M NEt4BF4 + AN solution; the inset: voltammogram on a 4 mm gold electrode at a scan rate of  2

1 0.1 V s
1. (B) Voltammograms corresponding to the redox process of O
(solid triangles), 12 kV s
1 (open triangles), 18 kV s
1 2 =O2 at 3 kV s
1  2

1 (open rectangles) and 30 kV s (open circles). (C) Voltammograms corresponding to the redox process of O
(open rectangles), 30 2 =O2 at 18 kV s kV s
1 (open circles), 48 kV s
1 (solid triangles), 60 kV s
1 (open triangles) and 100 kV s
1 (solid rectangles); the inset: Ipc vs. t plot. (D) DEp vs. t plot. Experimental data with 10 ns time lag correction (solid rectangles) and theoretical predictions (solid curve) for kad = 1.95 · 104s
1(a = 0.5) are all shown.

102

Z. Guo, X. Lin / Journal of Electroanalytical Chemistry 576 (2005) 95–103

pulse length of the double potential step chronoamperometry applied was as long as 40 ms. This kind of weak adsorption phenomenon could be observed only in Ôtruly transientÕ voltammetric experiments.

potential E. As the limiting current of steady-state voltammograms is I L ¼ 4nFDO rcO

ðA:4Þ

assuming that the oxidized form only is present initially in the solution. Then cO(0) and cR(0) can be expressed as

4. Conclusion  The electrochemical behavior of the O2 =O
2 and the
 2
O2 =O2 couple at a gold electrode in 0.9 M NEt4BF4 + AN solutions has been carefully studied. At low  couple is diffusion-controlled scan rates, the O2 =O
2 with a rate constant of 7.38 · 10
3 cm s
1. At high scan rates, it is adsorption-controlled with a rate constant of 5.21 · 104 s
1. At moderate scan rates, it is controlled by both diffusion and adsorption. The faradaic signal of the  2
O
2 =O2 couple was greatly distorted at low scan rates. However, if the scan rates were high enough, the disturbing reactions could be successfully Ôout-runÕ. Then, we could obtain kinetic information from the wellshaped voltammograms. In our experiments, the  was found, and the rate constant adsorption of O
2 was deduced as 1.95 · 104 s
1.

cO ð0Þ ¼ ðI L
IÞ=4nFDO r;

ðA:5Þ

cR ð0Þ ¼ I=4nFDR r:

ðA:6Þ

The dependence of kf and kb on potential is classical 0

k f ¼ k s exp½
anf ðE
E0 Þ;

ðA:7Þ 0

k b ¼ k s exp½ð1
aÞnf ðE
E0 Þ:

ðA:8Þ

Introducing Eqs. (A.5), (A.6), (A.7) and (A.8) into (A.3) leads to 0

ð4DO =pk s rÞ exp½anf ðE
E0 Þ 0

¼ ðI L
IÞ=I
ðDO =DR Þ exp½nf ðE
E0 Þ:

Using the Nernst equation and Eqs. (A.5), (A.6) we can obtain the current–potential dependence for a reversible reaction 0

ðDO =DR Þ exp½nf ðE
E0 Þ ¼ ðI L
I s Þ=I s :

Acknowledgement

ðA:9Þ

ðA:10Þ

Introducing Eq. (A.10) into (A.9), we obtain

This work was supported financially by the National Natural Science Foundation of China under Grant 20173054.

Appendix A Galus et al. [42] derived Eq. (A.1) to determine ks from steady-state microdisk voltammograms of a oxidation process assuming that the reduced form only is present initially in the solution: 0

0

ð4DO =pk s rÞ exp½anf ðE
E0 Þ ¼ ðI L
IÞ=I
ðI L
I s Þ=I s :

ðA:11Þ

Using Eq. (A.11), we can obtain ks from steady-state microdisk voltammograms of a reduction process with the oxidized form only present initially in the solution. Obviously, taking an oxidation current as positive would not change the results at all.

References

ð4DR =pk s rÞ exp½
ð1
aÞnf ðE
E0 Þ ¼ ðI L
IÞ=I
ðI L
I s Þ=I s :

ðA:1Þ

However, this equation could not be used directly in the case of a reduction process. So we derived an equation to solve it in our case according to method of Galus et al. In the case of a quasi-reversible system kf

O þ ne
¢ R: kb

ðA:2Þ

Taking a reduction current as positive, then the total current is I ¼ nFA½k f cO ð0Þ
k b cR ð0Þ;

ðA:3Þ

where cO(0) and cR(0) are concentrations at the electrode surface, A = pr2, while kf and kb are the cathodic and anodic heterogeneous rate constants for the cathodic and anodic process, respectively, at some given

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