Potential-Resolved Snapshot Impedance Spectroscopy for Exploring

Article pubs.acs.org/JPCC

Potential-Resolved Snapshot Impedance Spectroscopy for Exploring the Mechanism of a Complex Reaction Undergoing Parallel Pathways Byoung-Yong Chang*,† and Su-Moon Park*,‡ †

Department of Chemistry, Pukyong National University, 45 Yongso-ro, Nam-gu, Busan 608-739, Korea Interdisciplinary School of Green Energy, Ulsan National Institute of Science and Engineering, Ulsan 689-805, Korea

‡

ABSTRACT: Here we report an electrochemical snapshot approach with rate modulation, where impedance spectra are captured as snapshots on varied potentials. Using this method, we managed potential-resolved mechanistic analysis of a complex reaction composed of multiple and parallel steps. The snapshot impedance spectroscopy is based on transforming the E−I data in the time domain to those in the frequency domain, and the rate modulation of specific protonation steps is based on controlling the buffer capacity and pH of the solution. We applied this strategy to monitoring changes in impedances during p-benzoquinone reduction as a function of the scanned potential, which allowed us to elaborately analyze its chemical and electrochemical reaction steps in the frame of parallel electrochemical-chemical-electrochemical (ECE) and electrochemical-chemicaldisproportionation (ECD) mechanisms. As a result, we found that the second electron transfer of the reaction occurs either directly from the electrode or indirectly via disproportionation, and the relative contributions of each pathway depend on pH. While well-buffered solutions have mostly been used in studies in the past, poorly buffered solutions with appropriate buffer capacities are best taken advantage of for the purpose of controlling the protonation levels.

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INTRODUCTION A snapshot is a picture capturing an instantaneous moment. Recently, such a snapshot concept has been utilized in chemical studies for high-throughput data acquisitions and time-resolving analysis platforms. For example, rapid large-scale multianalyte detections were achieved,1 and dynamic molecular behaviors, structures, processes and related mechanisms were unveiled by capturing and analyzing instantaneous snapshots of transient chemical phenomena.2 While the snapshot technique can make many inroads into chemistry, it has had some difficulties in extending to electrochemistry due to chronological and potentio-dynamic behaviors of electrochemical processes. In this report, we developed electrochemical impedance spectroscopy (EIS) in the form of a snapshot method, and applied this technique for an intensive study on complex chemical reactions involving multichemical and electrochemical processes. Traditionally, an electrochemical technique aims at finding relationships between parameters such as the potential (E), the current (I), and the time (t), or vice versa. The electrochemical phenomena can then be understood by appropriately interpreting the E−I−t information.3 Use of multielectrode © 2012 American Chemical Society

arrays for parallel measurements can be used to certain extents for that purpose, but complex electrical connections and instrumentation would pose another problem. Therefore, the electrochemical approach to snapshot techniques has been hard to achieve on the basis of the E−I−t scaffold. However, transform of E−I data in the time-domain (t-domain) to another domain provides a new way to develop electrochemical snapshot techniques. Recently, snapshot voltammetry using a bipolar electrode has been reported, in which E−I in the t-domain was transformed to that in the space-domain.4 It was based on a potential window uniformly spread out in space rather than scanned in time, which led to the acquisition of a voltammogram within 1.5 s. The spatially distributed potential window was not only of analytical interest5 but used for electrochemical generation.6 Here, application of only a single potential, not multiple or Received: May 30, 2012 Revised: July 29, 2012 Published: July 31, 2012 18270

dx.doi.org/10.1021/jp305283z | J. Phys. Chem. C 2012, 116, 18270−18277

The Journal of Physical Chemistry C

Article

electrodes: a gold disk working (area = 0.020 cm2), a larger platinum gauze counter, and a Ag|AgCl (in sat’d KCl) reference electrodes. The working electrode was polished to a mirror finish with alumina powders of 14 μm sequentially down to 0.03 μm, followed by sonication in distilled water prior to each experiment. Cyclic voltammetric and snapshot impedance were simultaneously obtained by the staircase cyclic voltammetric-Fourier transform electrochemical impedance spectroscopy (SCVFTEIS) technique.8a,14b The potential was swept from 0.7 to −0.2 V at a speed corresponding to 50 mV/s, and chronoamperometric currents were obtained at every 5.0 mV for 100 ms at the rate of 50 k samples/s. Fourier transform converts the voltage and current profiles in the t-domain to impedances in the f-domain. Details about the procedure and calculation are described elsewhere.8a,14b,16

varying potentials, for a few seconds allowed potentiodynamic electrochemistry to be carried out. Another electrochemical snapshot technique is based on conversion of the E−I data in the t-domain to those in the frequency-domain ( f-domain) using Fourier transform. Thus, electrochemical signals recorded in the t-domain are represented as those in the f-domain such as FT voltammetry3b or FT-EIS (FT-electrochemical impedance spectroscopy).7 In this approach, a collection of ac waves of multiple frequencies is used as an excitation source for a short period and the response is then converted to that in the f-domain via Fourier transform. Consequently, an EIS spectrum is taken like a snapshot covering a wide frequency range. Applications of the technique include sequential snapshots taken in potentiodynamic8 and/or chronological9 modes, which provide rich information. In this work, we use the snapshot impedance method to probe transient stages of a fast complex electrochemical reaction while the rate of its reaction steps was properly modulated. We take p-benzoquinone (Q) reduction in aqueous media as the research target, which contains electron transfer and protonation steps as well as interactions between intermediate species. When Q undergoes reduction reactions on its way to QH2, different protonation levels are observed: 0protonation under an unbuffered condition,10 and 2-protonation in a well-buffered solution leading to the final destination, hydroquinone (QH2).11 The 1-protonation level is not stable because of the fast following second protonation.10b Nevertheless, in order to study the complex fastgoing mechanism, we managed to stabilize one of the intermediate stages using a poorly buffered solution of an appropriate buffer capacity somewhere between well- and unbuffered conditions. This way, we delayed the second protonation step long enough to intervene the reaction and take potentiodynamic snapshots of EIS spectra. We thus successfully combined two strategies to achieve our goal: modulation of a specific step of the complex reaction using a poorly buffered solution and acquisition of snapshot EIS spectra every 100 ms and 5.0 mV while the potential is scanned. Thus, the complex reaction was studied by stitching the snapshot impedance spectra along the potential scan.

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RESULTS AND DISCUSSION Transformation of Electrochemical Signals from tDomain to f-Domain. The key point of the snapshot approach is to transform the measured signals in the t-domain to another domain, where the time dependence is removed. The data are acquired in a short single time frame, and are decoded in the transformed domain.4 For example, the snapshot impedance spectroscopy transforms chronological voltage and current profiles to impedance spectra of a wide frequency range, which contain fundamental electrochemical information on the charge transfer (Rp), the mass transfer (σ), the electric double layer (Cd), and the solution resistance (Rs). Its theory is based on following equation (eq 1) describing the transient current produced upon a potential step (ΔV):14c 2 ⎤ ⎡⎛ ⎤ ⎡ 2 σ ⎞⎟ ⎥ 2σ −ΔV ⎢ ⎥ ⎢ ⎜ exp ⎜ erfc i(t ) = t t ⎢⎝ R p + R s ⎟⎠ ⎥ + R p + Rs R R ⎢ ⎦⎥ ⎣ p s ⎣ ⎦

+

■

−ΔVR p R s(R p + R s)

e−((R p+ R s)/(R pR sCd))t (1)

Although this is expressed as a function of time, it can be converted to a function of frequency by Fourier transform and visualized as an EIS spectrum (a Nyquist plot). Sequential snapshots recorded while time or the potential is varied provide chronological9 or potentio-dynamic8a,14b information of electrochemical reactions. Poorly Buffered Solutions. By definition, a well-buffered solution keeps its pH constant no matter how much proton is consumed or produced during a reaction. However, a practical buffer solution has a limited buffer capacity, β, of an expression,

EXPERIMENTAL SECTION All chemicals of ACS reagent grade were purchased from Aldrich (Milwaukee,WI) except for KCl and NaOH, which were from Samchun Chemicals (Seoul, South Korea). Buffer solutions of different pH values were separately prepared with different kinds and amounts of acids and bases depending on pH ranges according to procedures in references.12 The chemicals used here are HCl, citric acid, KH2PO4, boric acid, NaOH, Na2HPO4, and NaH2PO4. Their buffer capacities were made to be 0.0050 M/pH by controlling the total concentrations of acids and bases (see eq 2 below). Additional KCl was added to keep the ionic strength constant at 0.100 M. A homemade, fast rise potentiostat was the main device for the impedance measurements.13 An NI PCI-5412 Arbitrary Waveform Generator (National Instruments, Austin, TX) was used as a high speed potential step generator, which was fed into the potentiostat. The voltage and resulting current data were acquired through an NI PCI-5922 digitizer (National Instruments, Austin, TX) with 24-bit resolution, which was installed in a PC via a PCI interface.14 A Visual BASIC program was coded to control both the function generator and the data acquisition system.8a,15 The electrochemical cell had three

β=

⎛ K C K [H]+ ⎞ dn = 2.303⎜ W+ + [H]+ + buffer a + 2 ⎟ dpH (Ka + [H] ) ⎠ ⎝ [H] (2)

where n is the number of equivalents of a strong acid or base required to cause a unit change in pH per liter, Cbuffer is the total concentration of a weak acid and its conjugated base, KW is the water ionization constant, and Ka is the acid dissociation constant. For an electrochemical reaction involving protonation, its intermediate species captures protons and a sufficient β value is required to keep the pH constant; otherwise, the solution is not capable of providing enough protons, which affects the overall reaction. Thus, a well-buffered solution with a 18271



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high buffer capacity has been strongly recommended for studying pH effects. In this study, however, poorly buffered solutions of a low buffer capacity were employed to control the degree of protonation by allowing the first protonation but restricting the next. The buffer capacity used for reduction of 5.00 mM Q was 0.0050, which is equivalent to 5.0 mM of [H+] as the golden mean between well- and unbuffered solutions. When Q is reduced at the electrode surface and protonated to give QH•, the proton concentration would be depleted near the electrode surface. Consequently, the second protonation would be retarded until more protons are supplied from the bulk, during which period impedance snapshots could be taken without an interference of the subsequent protonation reaction. Rate Modulation by the Poorly Buffered Solution. Overall reduction of Q, that is, Q + 2H+ + 2e− ⇄ QH2, can be broken down to a series of electron transfer and protonation steps in a moderate pH medium (pH 3−9).11 Q + e− ⇄ Q−• kb

Q−• + H+ ⇄ QH• kf •

−

QH + e ⇄ QH + kd

at E10 ′

(3a)

K = k f /k b[H+]

(3b)

Figure 1. A series of cyclic voltammograms obtained from 5.00 mM Q reduction at a scan rate of 50 mV/s in poorly buffered solutions with their buffer capacity of 0.0050 at pH values ranging from 2.40 to 9.17. The peaks labeled (I), (II), and (III) represent the reduction of Q to QH−, the oxidation of QH− to Q, and the oxidation of QH2 to Q (see text). kb

A⇄B kf

−

QH− + H → QH 2

at

E20 ′

(3c)

with

K = k f /k b

(4)

Since the equilibrium during Q reduction involves protonation, K can be controlled by pH as long as the solution supplies enough protons to maintain the pH. Along with K, a kinetic parameter λ = (kf + kb)(RT/(Fυ)) is also important in describing the ECE mechanism of reactions 3a,3b, and 3c.3a,18 Here υ is the voltage scan rate. Replacing eq 3b by eq 4, both λ and K are redefined as λ = (kf + kb[H+])(RT/(Fυ)) and K = (kf/(kb[H+])) due to the role played by the proton reacting with Q•− to produce QH•.19 The thermodynamically and kinetically modulated reductions to show different CVs as seen in Figure 1. At pH = 9.17, the CV appears very reversible as the equilibrium constant, K = kf/(kb·[H+]), is very large while λ (= (kf + kb[H+])(RT/(Fυ))) is very small so that it would be located in the DO zone in the Savéant’s kinetic zone diagram.3a,18 This explains why the CV of 2e− reduction of Q in the buffered solution looks like that of 1e− reduction with slightly decreased peak currents and a broad peak potential separation. However, as [H+] increases and pH decreases, λ increases while K decreases so that CVs move from the DO zone to the DI zone, which is intermediate between one- and two-electron transfer zones.3a,18 In conclusion, the electrochemical kinetics of the intermediate stage is controlled by pH while the intermediate stage is stabilized by the appropriate buffer capacity of the poorly buffered solution. Two Pathways to QH−. What makes the benzoquinone reduction mechanism more complex is the second electron transfer step taking either of two different pathways: direct electron transfer from the electrode 3c and homogeneous electron transfer via disproportionation (designated as D) of two intermediate species described by

(3d)

While the pathway 3a through 3c was observed to be chemically reversible as an ECE reaction, reaction 3d renders the overall process irreversible, which causes the redox potential of the reaction to change. When protons are supplied abundantly in a well-buffered solution, Q reduction would go from reaction 3a all the way though 3d, but the protonation steps (3b and 3d) would be slowed down or hindered significantly when protons are not supplied due to their scarcity in an unbuffered solution. By formulating the poorly buffered solution with an appropriate buffer capacity, we selectively modulated the rates of two consecutive protonation steps: keeping the first protonation (reaction 3b), but restricting or delaying the second protonation (reaction 3d). When the first protonation is completed subsequent to the first faradaic reaction, the local pH on the electrode surface would be changed depending on the rate of the first protonation, which would determine the rate of the second protonation step. Even though we cannot quantitatively describe the degree of the restricted/delayed reaction rates of the second protonation due to the poorly buffered pH medium, its rate modulation is evidenced by the potential shift by pH changes, the modulation of reversal CV peaks shown in Figure 1, and the ratio of ECD/ECE reactions shown in Figure 4c. According to the relation, −(h/n) × 60 mV/pH, for an electrochemical reaction with n electrons and h protons involved, the poorly buffered solution having h = 1 and n = 2 will result in a potential shift of −30 mV/pH while the unbuffered and the well-buffered solutions show potential shifts of 0 and −60 mV/pH by h = 0 and 2, respectively. Similar potential shifts were reported with controlled amounts of protons in an aprotic medium.17 While the buffer capacity of the buffer solution controls the protonation levels, the pH value modulates the reaction rates relating to protonation. The equilibrium constant (K) of the first protonation 3b has a form of

QH• + Q•− → QH− + Q

(5)

Savéant et al. pointed out that the above reaction scheme should be considered as an alternative pathway, that is, ECD, instead of the straightforward ECE reaction when intermediate products favor disproportionation.18,19 Laviron also noted that the second electron transfer via disproportionation within the diffusion layer cannot be negligible.11,20 In this study, we found that both paths of electron transfer occur in parallel, that is, 18272



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reaction 3c is electrochemically measurable due to the direct electron transfer from the electrode, reaction 5 is not because no electron is transferred from the electrode. Nevertheless, the chemical change by reaction 5 can be measured due to change of chemical concentrations driving additional diffusional mass transfer effects, which is reflected on the Warburg impedance (ZW) in the equivalent circuit. Therefore, the ECE and ECD information on Q reduction can be explored by fitting the impedance data to the equivalent circuit and evaluating each element: Rs is the solution resistance, Rp1 and Rp2 are the polarization resistances for two electron transfer steps,14a Cd1 and Cd2 are the electrical double layer capacitances, and Zw is the Warburg or mass transfer impedance of all the electrochemically active molecules, whose reciprocal expression is the mass transfer admittance (Y0).10c Analysis of Mass Transfer. As described above, the mass transfer behavior can provide information on chemical reactions in terms of concentration changes because its mathematical expression is derived from the flux of chemicals toward the electrode surface due to the change in the concentration.10c,14b If there are two parallel pathways for the change in chemical concentrations, the total mass transfer admittance (Y0) made by all the redox couples is expressed as a sum of the two:

directly from the electrode surface and indirectly via the Droute. The ECE//ECD mechanism and contributions of each pathway depending on pH will be described using the impedance data. One more consideration should be noted about the number of electrons transferred (n) based on those mechanism frames. Even though n = 2 when Q is reduced to QH−, the apparent number of electrons transferred (napp) can be 2 or 1 depending on the pathway because they have different number of electrochemical reactions.3a When the reaction proceeds via ECE, the napp value would be 2, whereas it would be 1 if it proceeds via ECD. Impedance Spectra Taken during Potentiodynamic Scans. We obtained a series of impedance spectra from 0.70 V to −0.20 V versus Ag|AgCl at 5.0 mV and 100 ms intervals at various pH values. Figure 2 shows a set of typical data: (a) a

Y0 = Y1 + Y2 =

σi =

1 + 2 σ1

1 2 σ2

(7a)

⎧ DO, i ⎪ ⎡ nF ⎤ RT ⎨ exp⎢ − i (E − E1/2, i)⎥ 2 2 1/2 ⎪ ⎣ ⎦ RT 2 n F ADO Ci ⎩ DR, i +2+

DR, i DO, i

⎫ ⎡nF ⎤⎪ ⎬ exp⎢ i (E − E1/2, i)⎥⎪ ⎣ RT ⎦⎭

(7b)

where i denotes either of pathways (ECE or ECD), ni is the number of electrons transferred of pathway i, Ci is the concentration of the redox couple of pathway i, E1/2,i is the halfwave potential, and DO,i and DR,i are the diffusion coefficients of the oxidant and reductant involved in pathway i. Also, A is the electrode area, T is the temperature, R is the gas constant, and F is the faraday constant. In this particular system, the oxidants and reductants have molar masses of 108.1−110.1, which are different by less than 2% and, thus, DO,i/DR,i is taken to be 1.0. Figure 3 shows Y0 values, which were extracted from the impedance data,10c plotted as a function of potential in poorly buffered solutions at pH values 2.40 through 9.17 during cathodic (a) and anodic (b) scans. The E1/2 values, which were obtained from the potentials for peak admittances as shown by Figure 4a using eq 7, are plotted vs pH in panel c. A few points are noted from the plots. First, only one peak is seen during the cathodic scan (Figure 3a) while two peaks are observed during the reverse scan (Figure 3b). This indicates that one species is reduced during the cathodic scan while two different species are oxidized at two different potentials during the reverse anodic scan. Second, the slope of the E1/2−pH plot in Figure 3c is −30 mV/pH for cathodic peak I; however, the slopes are −30 and −60 mV/pH, respectively, for peaks II and III (data for peaks II and III not shown). These observations are in excellent agreement with CVs obtained with limited amounts of protons17 leading to the following assignments: peak I to reduction of Q to QH−, peak II to oxidation of QH− to Q, and peak III to oxidation of QH2 to Q, all by two electron transfers. In fact, the peak potentials of II and III were shown to shift by

Figure 2. (a) A current profile recorded for 0.100 s at E0′ (18 mV vs Ag|AgCl) at pH 5.56, (b) an impedance spectrum transformed from the voltage step (not shown) and the above current signal, and (c) equivalent circuit describing the impedance spectrum shown in panel b. The red solid line in panel b is the fitted curve for the data using the equivalent circuit shown in panel c.

current profile taken at 18 mV at pH 5.56, (b) its converted snapshot impedance spectrum, and (c) the equivalent circuit used for the spectral analysis. A depressed semicircle indicates that two one-electron transfers occur in series rather than a single-step two-electron process.21 Thus, the equivalent circuit has a serial combination of two RC loops as shown in Figure 2c. As mentioned above, the second electron transfer would take place via either of two pathways, reaction 3c or 5. While 18273



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protonation, resulting in an increase in pH within the diffusion layer and delaying further protonation. Also, different oxidation potentials observed for QH− and QH2 agree with those in the literature;17,22 the oxidation potential of QH− is more negative than that of QH2 because QH2 is thermodynamically more stable with its high pKa. Even though only a single peak is observed in buffered solutions during the cathodic scan, it is actually composed of the two successive electron transfer reactions with protonation intervening between them. When the 3c route is taken as the second electron transfer pathway, Q reduction becomes an overall two-electron process with n = 2. Alternatively, when the D route (reaction 5) is taken instead, it would be a one-electron process with n = 1 by reaction 3a because the ECD is one of the assorted EC mechanisms. For this reason, we resolved the mass transfer admittance into two processes by fitting the results with eq 7 and calculated the surface concentrations of the redox couples resulting from each process. Figure 4 shows (a) a typical mass transfer admittance plot obtained at pH 5.56 during the potential scan and (b) the concentrations involved in the ECE route (red square, pathway through reactions 3a, 3b, and c) and the ECD route (open circle, pathway through reactions 3a, 3b, and 5). With the diffusion coefficient of DO = 2.7 × 10−5 cm2/s,14a we find that the fraction of the disproportionation reaction increases as pH decreases while that of the direct electron transfer from the electrode decreases (Figure 4c). Concerning the reverse reaction of the disproportionation,23 its possibility becomes decreased in the presence of more protons. In conclusion, the homogeneous electron transfer via the disproportionation (ECD route) becomes more important in more acidic solutions in comparison to the direct electron transfer from the electrode (ECE route). Analysis of Charge Transfer Resistances. The charge transfer characteristics of Q reduction were studied by evaluating the polarization resistances. The Rp1 values were shown to undergo slight changes with the variation of the potential, generally much smaller than Rp2 values, and not dependent on pH. On the other hand, Rp2 shown with various colored lines in Figure 5 are dependent on the potential and pH. Thus, Rp1 must correspond to the first electron transfer to Q before protonation, while Rp2 corresponds to the second one (see Figure 2c), which is more important for the overall electron transfer process. Thus, it appears that the overall electron transfer reaction rate is determined by the second electron transfer as Rp1 and Rp2 are connected in series. The fact that Rp1 is not dependent on pH indicates that the first electron transfer has not much to do with the protonation step (reaction 3b), which is the same as the observation for Q reduction in unbuffered solutions, whose reduction potential is independent of pH.10b,22 The Rp2 values, however, change with pH as seen in Figure 5, and their potential profiles move in the negative direction as pH increases. The formal potential for the second electron transfer, E0′, at which Rp2 values and Warburg impedances have minimum values, shows a slope of −30 mV/pH as was the case for the E1/2 shift. Also, the Rp2 value increases as pH decreases and then levels off below pH ≈ 3.6. Exchange rate constants (k0) and charge transfer coefficients (α) of the second electron transfer from the electrode were obtained from the linear relationship between ln kf and the potential using the following equations:24

Figure 3. A series of the mass transfer admittances (Y0) obtained at different pH values ranging from 2.40 to 9.17 during: (a) cathodic and (b) anodic scans. The peaks (I), (II), and (III) also correspond to the CV peaks shown in Figure 1. (c) The E1/2 vs pH plot. The E1/2-values were taken from the admittance peak potentials at various pH values shown in panel a during the cathodic scan.

−30 and −60 mV/pH due to different deprotonation numbers, h = 1 and 2, respectively, as mentioned above. In the results, both QH− and QH2 were observed as reduction products in poorly buffered solutions whereas only QH2 was observed in well-buffered solutions. Due to the large pKa value of QH2 (around 10−12.5 depending on the conditions22), QH− is not expected to be observed in solutions of pH lower than 9. Nevertheless, it is observed in poorly buffered solutions because the proton concentration is rapidly depleted due to proton consumption upon the first 18274



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Figure 5. Charge-transfer resistances (Rp2) measured at different pH values are plotted along the potential.

ln k f = ln k0 − α

nF (E − E 0 ′) RT

(8b)

Here, Rp2 values in Figure 5 were used in the place of Rp in eq 8. A typical ln kf vs E plot is shown in Figure 6a for the impedance data for Q reduction at pH = 9.17, and the plot for the data obtained at pH = 2.40 is shown in Figure 6b over a wide potential range. Generally, the plots obtained at pHs lower than 7 had three linear regions, which are underpotential,

Figure 4. (a) Mass transfer admittances obtained (-■-) at pH 5.56, which were fitted with eq 7 and shown by the solid red line. The concentrations C1 and C2 obtained by fitting the data are surface concentrations of the redox pairs of the ECD and ECE routs, respectively. (b) C1 and C2 values obtained at different pHs of poorly buffered solutions. (c) The ratio of the ECE over ECD route plotted as a function of pH from C1 and C2 values.

kf =

⎡ ⎛ nF ⎞⎤ RT (E − E 0 ′)⎟⎥ ⎢⎣1 + exp⎜⎝ − ⎠⎦ RT 2n F AR pC*O 2 2

Figure 6. Plots of ln kf calculated from Rp2 values at (a) pH 9.17 and (b) pH 2.40 using eq 8. Although the kf values are different at different pH values, α values are constant at 0.41 at the region around the E0′, even at different pH values.

(8a)

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around E0′ and overpotential regions as can be seen in Figure 6b. In the around E0′ region, the transfer coefficient, α, was constant at around 0.41 regardless of pH but ln k0 decreased linearly as pH decreased (Figure 7). It is deduced that the

the stabilized (delayed) period. The rate modulation of the chemical reaction involving protonation was accomplished by controlling the buffer capacity and pH of the medium. Between zero- and two-protonation electrochemistry of Q reduction in unbuffered and well-buffered systems, the one-protonation state was established by slowing down the following protonation in a poorly buffered system, which is the golden mean of the two extreme buffer conditions. The application of snapshot impedances to the rate-modulated reaction system led to the successful study of the electrochemical reaction. Sequential impedance snapshots were taken every 100 ms and 5.0 mV during the potential scan; the mass and the charge transfer information obtained from independent single impedance snapshots was stitched on the potential axis. As a result, pbenzoquinone reduction was studied in the framework of the parallel ECE//ECD reaction mechanism at various pH values. The relative importance of the competing ECE and ECD reaction routes and the rate constant (k0) were evaluated along the pH. While the reduction was successfully studied in those frames, more complicated chemical phenomena such as the kinetics of the second protonation in a different medium, a quantitative study on the interface pH changed by faradaic reactions and so on can still be unveiled. For those in-depth studies, we are currently developing a new experimental technique in which both impedance and absorption spectroscopic experiments are run concurrently, and the results will be reported in due course. We have had experiences of resolving two different intermediate species observed during Q reduction employing two-dimensional analysis of spectroelectrochemical data and also sorting out various intermediate species employing a technique in which both electrochemical quartz crystal microbalance (EQCM) and spectroelectrochemical experiments were run simultaneously.25 In summary, two significant considerations accounted in terms of a chemical analysis technique in this study include the development of an electrochemical snapshot measurement technique to capture the momentous electrochemical processes and the paradigm shift about the buffer solution usage. Not only pH but the buffer capacity are worthy of being controlled for studies on the effects of protons on chemical reactions. A buffer solution of an appropriate buffer capacity should lead to the elucidation of reaction mechanisms of many other reactions, especially when protonation plays important roles.

Figure 7. The log k0 vs pH plot. The decreasing k0 with decreasing pH is presumably brought about by the increased importance of the disproportionation reaction as the second electron transfer process.

decreasing k0 results from the increased contribution of the disproportionation for the second electron transfer. In fact as shown in Figure 4c, the fraction of the second electron transfer to QH• from the electrode surface decreases at lower pH values in comparison to that via disproportionation, and this apparently raises Rp2 in Figure 5. This also causes the exchange rate constant to decrease at low pH values. As a result, k0 decreases along with the decrease in pH as shown in Figure 7. On the other hand, the Rp1 value is not affected by pH, as the first electron transfer occurs only directly from the electrode surface without protonation. An unexpected result was obtained for α. At pH = 2.40, for example, α was small (= 0.17) in the underpotential region, but it became large (= 0.94) in the high overpotential region (Figure 6b). The α value reflects the symmetry at the crossing point of the free energy surface potential profiles of the reactant and the product.3a Thus, the change in α-value for the second electron indicates that free energy surface potentials of both QH• and QH− do change depending on the potential regions at low pHvalues. While not clear, it is certain that the shapes of surface potential profiles of both species not only are different from each other but also change the depth as well as the widths of their vibrational motions by protonation depending on whether the potential is underpotential or overpotential. Finally, we see from Figure 7 that log k0 increases almost linearly as the pH increases, indicating that k0 decreases with an increase in the proton concentration. We attribute this to the relative decrease in the fraction of the direct two-electron transfer from the electrode surface as the proton concentration increases as was in Figure 4c. In other words, the apparent k0 value, rather than true k0, seems affected by the contribution of the disproportionation reaction during the second electron transfer.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]; [email protected]. Tel.: +8251-629-5597. Fax: +82-51-629-5583. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was supported by Basic Science Research Program (2011-0009714) and the WCU program (R31-2008000-20012-0) granted to UNIST through the National Research Foundation of Korea (NRF) funded by the Ministry of Education, Science and Technology.

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CONCLUSIONS The unstable intermediate stage during the p-benzoquinone reduction reaction was studied by two approaches: stabilization of the intermediate species by modulating the rates of intermediate steps and capture of impedance snapshots during

REFERENCES

(1) (a) Chow, K.-F.; Mavré, F.; Crooks, J. A.; Chang, B.-Y.; Crooks, R. M. J. Am. Chem. Soc. 2009, 131, 8364−8365. (b) Budach, W.; Neuschafer, D.; Wanke, C.; Chibout, S. D. Anal. Chem. 2003, 75 (11), 2571−2577.

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