Examination of Water Electrolysis and Oxygen Reduction As Self

Feb 22, 2011 - Suvi Lehtimäki , Anna Railanmaa , Jari Keskinen , Manu Kujala , Sampo Tuukkanen , Donald Lupo. Scientific Reports 2017 7, 46001 ...
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Examination of Water Electrolysis and Oxygen Reduction As SelfDischarge Mechanisms for Carbon-Based, Aqueous Electrolyte Electrochemical Capacitors Alicia M. Oickle and Heather A. Andreas* Department of Chemistry, Dalhousie University, 6274 Coburg Road, Halifax, NS, Canada B3H 4J3 ABSTRACT: Water electrolysis and oxygen reduction as possible self-discharge mechanisms for carbon-based, aqueous H2SO4 electrolyte electrochemical capacitors is examined through a comparison of the predicted and actual effects of varying the dissolved oxygen and hydrogen content on self-discharge. Water electrolysis, in the form of oxygen evolution, is not the self-discharge mechanism on the positive electrode, although, the selfdischarge profile is consistent with an activation-controlled Faradaic discharge or charge redistribution mechanism. The addition of hydrogen evidences no change in selfdischarge from the negative electrode, and the profile is consistent with a diffusioncontrolled mechanism, suggesting water electrolysis through hydrogen evolution is not the self-discharge mechanism. Oxygen reduction causes a large increase in self-discharge on the negative electrode which necessitates purging the cell of oxygen. As such, water electrolysis is likely not the cause of selfdischarge in carbon-based, aqueous electrolyte electrochemical capacitors but oxygen reduction is a cause of increased self-discharge on the negative electrode.

1. INTRODUCTION Electrochemical capacitors (ECs), also called supercapacitors, are energy storage devices that have attracted research attention in recent years. High cycle life and power capabilities make ECs an appealing option for certain applications (e.g., load leveling, regenerative braking, and camera flashes) compared to secondary batteries.1 However, ECs undergo a loss of energy during cell storage called self-discharge.1 This loss of energy over time is a limiting factor in the possible commercialization of ECs. Two possible reactions which may cause self-discharge in carbon electrode, aqueous electrolyte EC systems are examined here: water electrolysis and oxygen reduction. Electrolyte decomposition at potentials outside the region of electrolyte stability is a common problem in electrochemical systems using organic electrolytes, including organic-based ECs, and has been studied extensively since it results in decreased performance (e.g., refs 2-5 ). Water electrolysis has not been studied as a possible self-discharge mechanism for ECs using aqueous-based electrolytes. The possibility of water electrolysis occurring as a side reaction when aqueous-based ECs are charged near the maximum voltage has been previously mathematically modeled by Pillay and Newman to predict the diminished performance based on these side reactions.6 Pillay and Newman suggested, with reservations, that although the potentials of the EC electrodes stay within the thermodynamically stable potential window for aqueous electrolyte (0.00-1.23 V based on standard conditions at 298 K), water may still decompose to evolve oxygen or hydrogen, leading to a loss in electrode charge (loss in EC potential).6 Although Pillay and Newman were referring to water electrolysis as a side-reaction occurring concurrently with r 2011 American Chemical Society

EC charging, the suggested water electrolysis could also be a cause of self-discharge. If the necessary potentials for water electrolysis are reached, as would be expected for a charged EC, then the water electrolysis on the electrodes would cause the EC to spontaneously discharge (self-discharge). On the positive electrode, the evolution of O2 from H2O (reaction 1) would provide electrons which would discharge the positive electrode. On the negative electrode, the evolution of H2 from the protons in the acidic electrolyte (reaction 2) will cause the negative electrode to be discharged. Thus, the modeling by Pillay and Newman supports the possibility that water electrolysis may be a cause of self-discharge in aqueous ECs. 2H2 O f O2 þ 4Hþ þ 4e-

ðreaction 1Þ

2Hþ þ 2e- f H2

ðreaction 2Þ

Oxygen reduction via catalysts has been extensively studied for fuel cell research because O2 reduction occurs at the cathode of fuel cells.7-11 Pure carbon electrodes along with carbon electrodes with Pt have been used for O2 reduction to hydrogen peroxide (reaction 3),7 although they have shown decreased performance over time.7,11 Since oxygen reduction is possible on carbon electrodes at the potentials reached in ECs, this reaction is also a possible self-discharge reaction. If O2 reduction was occurring in the electrochemical setup, then oxygen would be Received: July 20, 2010 Revised: January 14, 2011 Published: February 22, 2011 4283

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removing electrons from the negatively charged electrode which would discharge the electrode and therefore this reaction must be studied to determine its effect on EC systems. O2 þ 2Hþ þ 2e- f H2 O2

ðreaction 3Þ

In the proposed self-discharge reactions above, there is an uptake/release of ions along with the transfer of electrons. In an electrochemical capacitor (or two-electrode system), the requirement of charge neutrality in the electrolyte ensures that the kinetics of the overall self-discharge for the system are limited by the slower of the self-discharge reactions. To study the selfdischarge reactions of the positive and negative EC electrodes separately, a three-electrode system must be used. This ensures that the kinetics of the reaction is governed by the self-discharge on the working electrode. Diagnostic kinetic models have been developed by Conway et al.,1,12,13 to aid in the identification of self-discharge mechanisms. The first self-discharge mechanism model, an activationcontrolled, nondiffusion-controlled Faradaic process, predicts a linear relationship between potential and log time. This profile would be the one expected for water electrolysis, as the high concentration of water in the electrolyte would guarantee that there are no diffusion limitations for the reactant. The second mechanism model predicts a linear relationship between potential and time1/2 for a diffusion-controlled Faradaic reaction.1,12,13 Although this relationship was derived for semi-infinite diffusion to a planar electrode, it was previously shown that the Conway equations can be applied to the high surface area carbons like those used in these studies since at low reactant concentrations the high surface area to volume ratio of the small pores results in rapid reactant depletion in the pores, meaning the pores deep within the carbon do not contribute to self-discharge and the electrode acts as a planar electrode.14 It is expected that oxygen reduction would exhibit a diffusion-controlled profile due to the low concentration of dissolved oxygen in the electrolyte. Plotting the self-discharge profiles using these criteria gives an insight as to the mechanism of self-discharge. In this work, water electrolysis and oxygen reduction has been experimentally studied as self-discharge mechanisms in aqueous electrolyte, carbon electrode ECs using the Nernst equation, along with Le Ch^atelier’s Principle, to predict the effect of varying the dissolved oxygen and hydrogen content on self-discharge from both the positive and negative electrodes. By identifying the mechanism of self-discharge through the Nernstian predictions and the Conway self-discharge mathematical models,1,12,13 a better understanding of, and possible decrease in, self-discharge may be realized.

2. EXPERIMENTAL SECTION 2.1. Electrodes and Electrochemical Cell. A glass threecompartment, three-electrode cell was used in all of the electrochemical experiments. The three-electrode setup was used, rather than a full-cell, two-electrode setup, to allow the selfdischarge mechanism of the positive and negative electrode to be determined independently, since the mechanism of each electrode is expected to be different. The carbon used in this study was Spectracarb 2225 carbon cloth (Engineered Fibers Technologies) with a BET surface area of ∼2500 m2 g-1. The working electrodes (∼10 mg of Spectracarb 2225 carbon cloth) were prepared in one of two ways: the first method involved holding the electrode in a Swagelok PFA tube fitting attached to a glass

tube housing the Pt wire used for the electrode connection. A silicone gasket was used to minimize electrolyte leakage into the glass tubing and to apply even pressure to ensure good electrical connection between the carbon, Pt mesh used as current collector and the Pt wire lead. Working electrodes were also prepared without Pt by using a 1:1 mixture of graphite carbon powder (Aldrich, 50 h is required to fully minimize the charge redistribution effect.19 Since a 50-h hold step is prohibitively long for these experiments, and since it is expected that this long of a hold step may change the Faradaic reactions occurring on the surface (since during the hold step, the Faradaic reactions causing selfdischarge would be expected to be occurring, likely building up the concentration of the product and changing the conditions near the electrode surface), a shorter hold step was used in this work. It was previously shown that a 30 or 60 min hold time 4284

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greatly reduces the charge redistribution14,19 and therefore both of these hold times were examined here. Since both holding times resulted in the same trends, and these trends agree with experiments where no holding step was used, only the 30 min holding step self-discharge data was shown in this paper. The choice of the 30 min holding step was to minimize charge redistribution, which also has a linear relationship when plotted versus log t, and still allow the Faradaic reaction which may be discharging the electrode to be visible.20

3. RESULTS AND DISCUSSION 3.1. Water Electrolysis. 3.1.1. Self-Discharge Profiles from Electrodes Charged to 1.0 V. Predictions using Le Ch^atelier’s

Principle and the Nernst equation suggest that if water electrolysis, in the form of water decomposing to O2 (reaction 1) on the electrode charged to 1.0 V, was the self-discharge reaction the addition of oxygen content to the electrolyte would have a positive shift in the reaction potential (farther out of the potential window), resulting in a decrease in self-discharge from 1.0 V, evidenced as less potential change during the self-discharge measurement. Similarly, the removal of O2 from the electrolyte (by sparging with N2) should result in a negative shift in reaction potential and a concomitant increase in self-discharge. As seen in Figure 1a, with the addition of oxygen by bubbling O2 through the cell, there is no significant effect on self-discharge. In addition, the removal of oxygen from the electrolyte, by bubbling N2 through the cell, results in no change in self-discharge within experimental error. Since the self-discharge profiles do not follow the trend predicted based on the amount of dissolved oxygen in the electrolyte, this suggests that water electrolysis is not the self-discharge mechanism in the carbon electrode, aqueous electrolyte ECs. The evidence against water electrolysis can be strongly supported by calculations of the O2 content required to shift the reaction potential into the voltage window accessed by the electrode. Assuming the electrolyte is not deoxygenated, the dissolved oxygen content can be calculated using Henry’s Law and a typical atmospheric oxygen content of 0.21 atm. Combining this dissolved oxygen concentration with a 1.0 M H2SO4 into the Nernst equation would result in only a -50 mV shift in the potential to 1.18 V. The quantity of oxygen necessary to shift the equilibrium potential even to 1.0 V (the very upper limits of the potential window) would be 2.8  10-16 mol L-1, a trace amount unlikely to be reached under normal conditions. Indeed, given the large overpotentials which would be required for the reaction to occur, a concentration which is even smaller than the one calculated here would be required in order to push the potential down to where the reaction would occur to any extent. Dissolved oxygen probe measurements showed the quantity of O2 in an N2 saturated solution to be 0.10 atm O2 (1.3  10-4 mol L-1), which is significantly different from the value necessary to move the water oxidation reaction into the working potential window. To further elucidate the kinetics of the self-discharge reaction occurring on the positive electrode, the self-discharge profiles were plotted versus log t as well as t1/2. A water electrolysis mechanism is expected to be activation-controlled, rather than diffusion-controlled. With O2, N2, and no bubbling, there is a linear region, after an initial plateau, in the self-discharge profile when plotted versus log t (Figure 1b), which is consistent with an activation-controlled self-discharge mechanism. However, it is also consistent with charge redistribution14,19 which may not have been completely removed in the 30 min hold. That

Figure 1. Self-discharge profiles of Spectracarb 2225 carbon cloth, in 1.0 M H2SO4, charged to 1.0 V at 1 mV s-1 with a 30 min holding step. Solution not bubbled (red dashes); Saturated with N2 (green dashes and dots); Saturated with O2 (blue line).

the self-discharge mechanism is not diffusion-controlled is further confirmed by the lack of a linear region in the selfdischarge profile when plotted versus t1/2 (Figure 1c). From this data, it is shown that while the self-discharge mechanism may be activation-controlled, the lack of trend with varying dissolved O2 concentration argues against self-discharge at the positive electrode being caused by water electrolysis, in the form of O2 evolution. Rather the self-discharge is either another, as yet unidentified, activation-controlled Faradaic reaction or is governed chiefly by charge redistribution. 4285

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The Journal of Physical Chemistry C 3.1.2. Self-Discharge Profiles from Electrodes Charged to 0.0 V. When the working electrode is charged to 0.0 V in 1.0 M H2SO4, the water electrolysis reaction would be in the form of protons being reduced to H2. Under typical conditions (no bubbling of any gas), the dissolved H2 concentration in the electrolyte solutions is 6.2  10-7 mol L-1, using Henry’s Law and atmospheric H2 content. Using the Nernst equation, this H2 concentration in 1.0 M acid results in a positive potential shift of 0.18 V, which is a shift into the available voltage window, suggesting that water electrolysis is a reasonable self-discharge reaction. The addition of H2 would be predicted, using the Nernst equation, to cause the reaction potential to shift negatively, causing a decrease in self-discharge, and the removal of H2, through N2 sparging, should cause an increase in self-discharge. As can be seen in Figure 2a, the self-discharge profile of the negative electrode with the addition of H2 was not significantly changed (within experimental error) from profiles where the electrolyte was N2 saturated to remove dissolved H2. The small variation seen is also inconsistent with the predictions based on water electrolysis as the self-discharge mechanism. The self-discharge profile for the nonbubbled system is surprisingly large, and is out-of-line with the prediction of water electrolysis. The results for this nonbubbled system will be discussed again in section 3.3, but, in general, these results argue against water electrolysis as the self-discharge mechanism on the negative electrode. The kinetic relationships of the self-discharge profiles, when compared to the Conway models, also provide strong evidence that water electrolysis is not the dominant self-discharge mechanism on electrodes charged to 0.0 V. That the self-discharge mechanism is not activation-controlled is suggested by the nonlinear relationship when plotted versus log t (Figure 2b). The self-discharge profiles are, rather, consistent with a diffusioncontrolled self-discharge mechanism, as can be seen with a linear region in Figure 2c. Water electrolysis should not provide a diffusion-controlled profile, as the electrolyte does not need to diffuse to the surface to react, and should, instead result in an activation-controlled mechanism. Therefore, the presence of a diffusion-controlled profile also strongly negates the possibility of water electrolysis being the self-discharge mechanism on the negative electrode in carbon electrode, aqueous electrolyte ECs. 3.2. Oxygen Reduction on the 0.0 V Electrode. The presence of oxygen in the EC electrolyte provides another possible self-discharge reaction in the form of oxygen reduction to water or hydrogen peroxide. Hydrogen peroxide is the expected product for the system under study, as is typical with carbon electrodes.7 Self-discharge profiles were collected from 0.0 V with varying concentrations of dissolved oxygen in the electrolyte to provide insight into a possible oxygen reduction self-discharge mechanism. Oxygen reduction as the self-discharge reaction would be consistent with the larger degree of self-discharge exhibited by the nonbubbled system (cf., the N2-saturated and H2-saturated systems) in Figure 2. Bubbling either N2 or H2 would be expected to remove O2, preventing O2 reduction and therefore resulting in less of a potential change if oxygen is the self-discharge mechanism. Additionally, the low concentration of O2 present in either the nonbubbled or N2/H2 bubbled cases would be expected to result in a diffusion-controlled selfdischarge reaction, as indeed is seen in the linear V vs t1/2 profile in Figure 2c at longer times. Near the beginning of the self-discharge profiles, the system is likely under mixed control, where the reaction rate may depend on activation-control (with an exchange current density similar to the 10-4 A cm-2 rate for graphite or carbon nanotubes21)

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Figure 2. Self-discharge profiles of Spectracarb 2225 carbon cloth, in 1.0 M H2SO4, charged to 0.0 V at 1 mV s-1 with a 30 min holding step. Solutions not bubbled (red dashes); Saturated with N2 (green dashes and dots); Saturated with H2 (black dashes and double-dots); Saturated with O2 (blue line).

charge redistribution and/or transport-control. At longer times, the self-discharge current, I, (calculated by the following: dV ð1Þ I ¼ -C dt where C is the capacitance1,12,13) falls below that of the exchange current density suggesting that diffusion is limiting self-discharge. Thus, the results seen in Figure 2 are consistent with oxygen reduction as the self-discharge mechanism on the negative electrode, with diffusion as the rate determining step at longer self-discharge times. 4286

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Further support for oxygen reduction being the self-discharge mechanism was seen when O2 was bubbled through the electrolyte prior to (24 h) and during self-discharge. The increase in dissolved oxygen content resulted in a significant increase in the self-discharge (a much larger potential change over time), as seen in Figure 2, and suggests that minimization of the O2 content in the electrolyte is necessary when the electrode is negatively charged. The concentration of O2 near the electrode or in the pores which is required to cause the degree of self-discharge seen in the nonbubbled system in Figure 2, can be estimated from the loss in potential (∼0.24 V), the capacitance of the carbon in 1.0 M H2SO4 at 0.0 V (155 F g-1), Faraday’s constant (96 485 C per mole of electrons), the pore volume of the carbon (1.2 mL g-1) and 2 electrons being used for the reduction of one O2 molecule to hydrogen peroxide. This estimation suggests that a dissolved oxygen concentration of 0.16 mol L-1 in the pores is required to cause the degree of self-discharge evident in these systems. The oxygen content of a nonbubbled electrolyte is ∼1.6  10-4 mol L-1 (as measured by the dissolved oxygen probe). Thus, the amount of oxygen in the pores is insufficient to fully account for the degree of self-discharge seen in Figure 2 and the discharge would require O2 from the bulk electrolyte to diffuse to the electrode surface, and may account for the t1/2 profile evident in this system. The linearity of the self-discharge profile when plotted versus t1/2 (as seen in Figure 2) argues for the electrolyte in the pores having been depleted of dissolved oxygen, as this t1/2 profile was modeled for semi-infinite diffusion to a planar electrode.1,12,13 In contrast, if the limiting reaction was diffusion in the pores of the electrode, then a much more complex relationship between the flux to the surface and time is expected. For diffusion in a cylindrical pore, the flux (F) and time would be related through the equation modeled by Crank22 for long times: " # 2D 1 γ ðC1 - C0 Þ F ¼ - ::: a lnð4TÞ - 2γ flnð4TÞ - 2γg2 ð2Þ where D is the diffusion coefficient, a is the radius of a cylinder, C1 is the concentration at time 1, C0 is the bulk concentration, T = Dt/a2, and γ is Euler’s constant (0.57722). Since the flux of material is directly related to the current, and the current discharges the electrode, thereby changing the electrode potential, the self-discharge potentials would also be expected to have a similar relationship with time if there is diffusion in a cylindrical pore. Pore narrowing and pore size polydispersity are also expected in this system and will also cause a more complex relationship between voltage and time. Thus, the clear linearity in self-discharge potential with square-root time argues that the diffusion-control Faradaic reaction controlling self-discharge does not occur inside the pores to any great extent. This suggests the O2 reactant is depleted in the pores during the electrode charging process or early on in the self-discharge. The electrodes in this work would likely result in this planar-type of diffusion because of the small pores in this carbon, resulting in a high surface area to volume ratio in the pores. This allows for the pores to be depleted of oxygen rapidly, and may minimize the active surface area for Faradaic self-discharge based on diffusion. As the charge on the external surface of the electrode is removed through the Faradaic self-discharge process, the charge

Figure 3. Comparison of current profiles of Spectracarb 2225 carbon cloth, in 1.0 M H2SO4, charged to 0.0 V at 1 mV s-1 during a 30 min holding step. Inset is a magnification at low current. Solutions not bubbled (red dashes); Saturated with N2 (green dashes and dots); Saturated with H2 (black dashes and double-dots); Saturated with O2 (blue line).

deeper in the pores is likely fed up to the surface through charge redistribution.19,20 Through this process, the whole electrode surface may eventually be discharged. To examine if the pores may be depleted during the ramp and subsequent hold step in charging of the electrode, the current recorded during the holding step is shown in Figure 3. It can be seen that the current was higher for the nonbubbled situation versus the system when bubbled with either H2 or N2, as is consistent with a higher concentration of O2 in the electrolyte. (Note the glitch seen ∼1300 s in the H2 and N2 systems is due to instrumentation and can be ignored for this comparison). The curve for the nonbubbled system is, however, approaching the curves for the system where O2 has been removed, suggesting that the O2 in the pores may be in the process of being depleted during the holding step. The equilibrium current never reaches the level of the H2 and N2 currents, always remaining slightly higher. This slightly higher current is likely due to O2 diffusing from the bulk and reacting on the face of the electrode. Again, this is consistent with the linear potential drop with t1/2 (seen in Figure 2c), which the Conway model suggests is due to a diffusion-controlled process1,12,13 and suggests that the external surface is where the majority of the self-discharge reaction occurs. From the current profiles seen in Figure 3 it can be seen that the current of the O2-bubbled electrolyte has a larger cathodic current during the holding step which is consistent with the entire O2 content in the pores not reacting during the 30 min holding step. Had the entire O2 content been used up during the holding step, the current profile would have been similar to those for N2-bubbled or H2-bubbled systems. The higher current seen in Figure 3 suggests more O2 reacting on the surface, arising from a higher oxygen content of the electrolyte, and therefore the higher concentration of the O2 in the pores. This higher concentration, at least initially, may be sufficient to lead to a mixed activation- and diffusion-control situation, explaining why the curve in Figure 2c for the O2-bubbled system is not linear with t1/2 (diffusion-controlled) at short times. As the O2 in the pores is used up during self-discharge, the system may become more diffusioncontrolled as the reaction becomes limited by transport of the O2 from the bulk electrolyte to the surface of the electrode, and 4287

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The Journal of Physical Chemistry C therefore resulting in the diffusion-controlled profile at longer times, as is seen in Figure 2c. The N2-bubbled and H2-bubbled systems exhibit considerably less self-discharge than the nonbubbled and O2-bubbled systems, but their self-discharge is not insignificant. It is possible that enough oxygen remains in the electrolyte that it is responsible for this selfdischarge. Alternately, it is another diffusion-controlled self-discharge reaction all together and experiments are continuing to test these two hypotheses. Nevertheless, O2 reduction has been identified as a main cause of self-discharge, suggesting that removal of O2 from an electrolyte will minimize self-discharge in ECs.

4. CONCLUSIONS Studies into possible self-discharge mechanisms in symmetric carbon electrode, aqueous electrolyte ECs were carried out. Predictions were made using the Nernst equation and Le Ch^atelier’s Principle, and carried out experimentally by varying the dissolved oxygen and hydrogen content, individually. Comparisons of self-discharge profiles showed that self-discharge is not caused by water electrolysis. The 1.0 V (positive) electrode did not show the expected decrease in self-discharge upon the addition of O2, indicating that water being oxidized to O2 was not the mechanism of self-discharge. With regard to the 0.0 V (negative) electrodes, the self-discharge profile indicated a diffusion-controlled mechanism, which when coupled with the absence of any relationship between H2 content and degree of self-discharge suggests that the water electrolysis reaction of protons being reduced to water is not the mechanism of self-discharge. In varying the dissolved oxygen content of the electrolyte, the self-discharge profile of the negative electrode was shown to exhibit significantly larger potential change on open-circuit (more selfdischarge) with increasing oxygen concentration. This result is consistent with oxygen reduction to hydrogen peroxide as being a main cause of self-discharge on the negative electrode. Additionally, although sparging of the electrolyte with N2 removes O2 and lessens the amount of self-discharge experienced, enough oxygen remains to allow for some degree of self-discharge. The kinetics of the oxygen reduction, while slow, are sufficient to account for the degree of selfdischarge seen in these systems. A self-discharge profile that is linear in square-root of time suggests that the electrolyte in the pores of the carbon has been depleted of dissolved oxygen, and that the selfdischarge is limited by diffusion of oxygen to the external face of the electrode, and that the pores do not directly play a role in the Faradaic self-discharge process. It is likely, however, that the pores may feed their charge up to the surface through charge redistribution, allowing the whole electrode to eventually be discharged through the external electrode surface area.

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’ AUTHOR INFORMATION Corresponding Author

*Tel: 902-494-4505. Fax: 902-494-1310. E-mail: heather.andreas@ dal.ca.

’ ACKNOWLEDGMENT The authors gratefully acknowledge the Natural Sciences and Engineering Research Council (NSERC) for funding of this work. ’ REFERENCES (1) Conway, B. E. Electrochemical Supercapacitors; Kluwer Academic: New York, 1999. 4288

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