Flow Battery Electroanalysis: Hydrodynamic Voltammetry of Aqueous

Subscriber access provided by University of South Dakota

Article

Flow Battery Electroanalysis: Hydrodynamic Voltammetry of Aqueous Fe(III/II) Redox Couples at Polycrystalline Pt and Au Tejal Sawant, and James R. McKone ACS Appl. Energy Mater., Just Accepted Manuscript • DOI: 10.1021/acsaem.8b00859 • Publication Date (Web): 04 Sep 2018 Downloaded from http://pubs.acs.org on September 8, 2018

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Applied Energy Materials

Flow Battery Electroanalysis: Hydrodynamic Voltammetry of Aqueous Fe(III/II) Redox Couples at Polycrystalline Pt and Au Tejal V. Sawant and James R. McKone∗ Department of Chemical and Petroleum Engineering, Swanson School of Engineering, University of Pittsburgh, Pittsburgh, PA 15261, USA E-mail: [email protected]

Abstract The redox flow battery (RFB) is a promising technology for large-scale electrochemical energy storage, but research progress has been hampered by conflicting reports of electrontransfer rates even for well-established battery chemistries. To address this challenge, we are working to deploy established electroanalytical techniques for precise characterization of RFB reaction kinetics. We studied Fe3+/2+ redox chemistry using rotating-disk electrode voltammetry with polycrystalline Pt and Au working electrodes as a model of an Fe/Cr RFB positive electrolyte. Our measurements yielded exchange current densities of 3.7 ± 0.5 and 1.3 ± 0.2 mA cm−2 for Pt and Au, respectively, in electrolytes containing 5 mM each of Fe3+ and Fe2+ . Both the variability and relative sluggishness of these rates are clear evidence that innersphere (catalytic) processes are important even in the 1-electron redox chemistry of Fe aquo complexes. Increasing the Fe concentration by 100-fold gave exchange current densities at Pt that were only ∼15-fold higher, suggesting that the reaction is not first-order in Fe or that the predominant mechanism changes as electrolyte concentration is increased. These results

1

ACS Paragon Plus Environment

ACS Applied Energy Materials 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

motivate further studies of RFB electrocatalysis under conditions that prioritize both analytical precision and device-level applicability.

Keywords Energy conversion, Redox Flow Battery, RFB, Kinetics, Iron, Fe, Platinum, Pt, Gold, Au, Electroanalysis, Rotating disk electrode voltammetry, RDE

Introduction Energy demand worldwide is massive, and our current energy system is widely seen as environmentally unsustainable. Renewable energy technologies, particularly wind and solar electricity, are highly attractive methods to decarbonize the global energy supply. 1–3 However, these types of renewable resources are spatially and temporally intermittent, which would make it difficult to maintain a stable electric grid based predominantly on wind and solar power. 4–7 Grid-scale electrochemical energy storage can help mitigate this problem by providing large quantities of renewable electricity when and where it is needed. 8–10 The redox flow battery (RFB) is one promising method for grid-scale electrochemical energy storage. 11–14 RFBs combine elements of conventional secondary batteries (containing solidphase active materials) and fuel cells. Like conventional batteries, RFBs operate via reversible redox reactions between two sets of electroactive materials, such as transition metal complexes or redox-active organics, which are stored within the battery assembly itself. However, in RFBs the electroactive materials are dissolved or suspended in liquid media to form positive and negative electrolytes, which flow through a charge-discharge stack to interconvert between electrical and chemical energy in a process resembling the operation of a fuel cell. RFB electrolytes can be stored in tanks of arbitrary size; these tanks and the associated charge/discharge stacks can be sized

2


Page 2 of 31

Page 3 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


independently, which provides flexibility in the relationship between energy and power. These characteristics make RFBs especially attractive for grid-scale energy storage if suitable component and system-level cost targets can be met. 15 RFBs have been the subject of considerable attention in laboratory studies and in commercial development, as exemplified by the extensive review literature. 16–24 Early work supported by the US National Aeronautics and Space Administration (NASA) focused on Fe and Cr salts dissolved in aqueous acidic solutions. 25–27 The all-vanadium RFB, which uses V3+/2+ and V5+/4+ redox couples as the respective negative and positive active materials, has also been studied extensively. 28–36 Considerable work is now focused on the development of new electrolytes, resulting in a growing library of active materials for aqueous RFBs. 37–42 There is also increasing work on nonaqueous RFB active materials, particularly electroactive organics and organometallics, which offer the advantage of high cell voltage due to the extended stability window of organic solvents. 43–47 A major challenge for the design and implementation of RFBs involves minimizing efficiency losses associated with sluggish electron transfer from the positive and negative electrodes to the corresponding electrolyte species. For example, the aqueous Fe3+/2+ redox couple serves as the positive electrolyte in Fe/Cr RFBs, and its electron transfer kinetics have been measured using various methods, including hydrodynamic voltammetry, 48–56 AC impedance, 57 faradaic rectification, 58 and electrochemical pulse techniques. 59,60 These studies reported values of the interfacial electron-transfer rate constant, k 0 , that vary over several orders of magnitude even for relatively well-understood electrode materials like polycrystalline Pt and Au (see Supporting Information for tabulated values). This variability is highly problematic for RFB design, as illustrated in Figure 1 in which we have simulated the overpotential performance of a hypothetical RFB positive electrode at the lower and upper limits of reported k 0 values for Fe3+/2+ at polycrystalline Au. Notably, these limits span the range over which kinetic limitations would begin to contribute significantly to efficiency losses in a functional device. Moreover, this level of variability is not restricted to iron redox chemistry; vanadium RFB electrolytes have also been the subject of debate over which redox couple exhibits more sluggish electron-transfer kinetics. 61–67 3



Figure 1: Simulated performance of a Fe3+/2+ RFB positive electrolyte in terms of overpotential versus state of charge under galvanostatic operation. The heterogeneous electron transfer rate constant was varied between k 0 = 10−2 and 10−5 cm s−1 , corresponding to the outer bounds of reported values at Au electrodes. Complete simulation details are described in the Supporting Information. We are working to accurately characterize the interfacial chemistry of RFB electrodeelectrolyte combinations using well-established electroanalytical tools in the interest of resolving ambiguity regarding electrode kinetics and ultimately improving our ability to design efficient devices. Thus, we have deployed an analytical protocol for kinetics measurements based on rotating disk electrode (RDE) voltammetry, which was modeled after analogous techniques for characterization of fuel cell and water electrolysis catalysts. 68–71 Our approach emphasizes rigorous electrode surface preparation, which we have found to be extremely important to obtain reproducible kinetics in RFB-mimicking electrolytes. We report here a detailed description and critical analysis of our experimental protocol as it was applied to the redox chemistry of Fe3+/2+ using polycrystalline Pt and Au electrodes. We found exchange current densities (j0 ) of the Fe3+/2+ redox reaction to be 3.7 ±0.5 and 1.3 ±0.2 mA cm−2 on Pt and Au electrodes, respectively, in electrolytes containing 5 mM concentrations each of Fe3+ and Fe2+ in 0.5 M HCl. Importantly, the electrode surfaces needed to be repeatedly refreshed using electrochemical cleaning methods to obtain reproducible results. We also applied our analytical approach to electrolytes with 1 M total Fe concentration to more closely

4


Page 4 of 31

Page 5 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


resemble a practical RFB. Even after compensating for series resistance, we measured j0 = 60 mA cm−2 at polycrystalline Pt, which was ∼6-fold slower than would be predicted for a 1-electron transfer reaction whose mechanism is first-order in Fe. Additional experiments to probe the effects of anion adsorbates and preliminary results on carbon working electrodes further illustrate that RFB interfacial electron transfer dynamics are complex and merit careful scrutiny using the tools of applied electroanalysis.

Experimental Hydrochloric acid (Certified ACS plus), sulfuric acid (trace metal grade), FeCl2 (tetrahydrate salt, 98%) and FeCl3 (hexahydrate salt, 97%), were obtained from Fisher Scientific and were used without further purification. All solutions were prepared using deionized water with resistivity ≥18.2 MΩ·cm (Millipore, Milli-Q Advantage A10). Commercial acidic and alkaline detergents (Citranox and Alconox, respectively) were obtained from W.W. Grainger Inc. Alumina powders and polishing pads (Micropad) were obtained from Pace Technologies. Graphite electrodes were spectroscopic grade with a diameter 1/4 inch and porosity of 16.5 % and were obtained from Electron Microscopy Sciences. Reference electrodes were Ag/AgCl gel-type electrodes with 3 N NaCl fill solution and were obtained from Fisher Scientific. Ultra-high purity H2 (g) (99.999%) and zero grade N2 (g) (99.998%) were obtained from Matheson. Figure 2 depicts the experimental setup for these measurements. The RDE apparatus was a Pine MSR rotator equipped with ChangeDisk polycrystalline Pt and Au electrodes that were 5 mm in diameter. All electrochemical measurements were performed using a Gamry Interface 1000E potentiostat. The electrochemical cell used was a 100 mL glass chamber equipped with a tight fitting polytetrafluoroethylene (PTFE) cap into which holes were drilled to introduce electrodes and gas flow tubing. An electrolyte bridge was used to separate the reference electrode from the working chamber.

5



Figure 2: Schematic and photograph (taken by the authors) of the RDE experimental setup. Components labeled in the schematic are as follows: (a) glass cell, (b) teflon cap, (c) nitrogen purge tube, (d) vent, (e) RDE motor, (f) RDE shaft, (g) electrolyte solution, (h) working electrode, (i) counter electrode, (j) stopcock, (k) electrolyte bridge, (l) Ag/AgCl reference electrode and (m) reference electrode compartment. A representative experimental procedure for executing RDE measurements with polycrystalline Pt and Au working electrodes is as follows. First, the electrochemical cell and the reference electrode chamber were cleaned in 1 wt% aqueous Citranox solution by boiling for 30 mins followed by rinsing 10 times with pure water. The same treatment was then repeated using 1 wt% aqueous Alconox solution. Meanwhile, 50 mL of 0.5 M sulfuric acid solution was prepared, and the equilibrium potential of the Ag/AgCl reference electrode was calibrated against a reversible hydrogen electrode (RHE) with this solution using a clean Pt button electrode. A representative result for this calibration was EAg/AgCl = 0.233 V vs. RHE, which is in close agreement with the predicted value of 0.227 V vs. RHE (0.209 V vs. NHE). Next, a stock solution of 150 mL of 0.5 M HCl was prepared. 50 mL of this solution was added to a 50 mL polyethylene centrifugation tube and appropriate quantities of FeCl2 · 4 H2 O and FeCl3 · 6 H2 O were added to produce a solution of 5 mM FeCl2 and 5mM FeCl3 . A few mL of this solution were used to rinse the electrochemical cell and discarded; the balance was then added to 6


Page 6 of 31

Page 7 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


the cell for data collection. For the preparation of the working electrode, 3 separate pieces of polishing pad were impregnated with 5, 1 and 0.05 µm alumina slurries in water, respectively. The working electrode was first polished using the 5 µm alumina slurry in a circular fashion with successive sets of 20 clockwise and counter-clockwise rotations for a total of 1 minute. It was then rinsed and sonicated to remove alumina residue by submersing in a water-filled 20 mL glass vial that was placed in a sonicator bath (Branson M1800) for 30 seconds. This polish–rinse–sonicate procedure was repeated using the 1 and 0.05 µm alumina slurries. The electrode was finally rinsed thoroughly with water, attached to the electrode rotator assembly, and lowered into the electrochemical cell without being allowed to dry. We separately characterized the surface morphology of a Pt electrode prepared in this manner, which exhibited residual submicron surface roughness after polishing. Complete details are included in the Supporting Information. The graphite counter electrode was cleaned using gentle abrasion with a paper wipe and placed in the same chamber as the working electrode. To connect the reference electrode, a separate 20 mL glass vial was connected to the main cell using an electrolyte bridge, which consisted of a length of polymer tubing (Flexelene, Cole Parmer) coupled to a PVDF stopcock. To fill the vial with electrolyte from the main cell, first a syringe was used to draw an initial quantity of solution from the main chamber into the tubing, after which the stopcock was closed to hold the solution in place. Then the opposite end of the tubing was placed into the vial, which was positioned at a lower height compared to the main cell to allow the flow of electrolyte by gravity. Upon opening the stopcock, the electrolyte began flowing into the vial from the cell, then the vial was raised and the liquid levels were allowed to equilibrate so that the reference chamber contained ∼5 mL of electrolyte. The calibrated Ag/AgCl reference electrode was placed in the vial to complete the cell setup. After cell assembly and prior to experimentation, the working electrode chamber was sparged with N2 (g) by bubbling through the electrolyte using Tygon tubing coupled to a glass pipette as a sparge tube. The sparging proceeded for at least 5 minutes, which was independently determined to be sufficient amount of time to remove O2 (aq) as described in the 7



Page 8 of 31

Supporting Information. Finally, the working electrode was cleaned electrochemically by cycling in the potential range of -0.25 V to 1.25 V versus Ag/AgCl at 100 mV s−1 for 20 cycles. Experimental data collection was performed as a repeating sequence of clean and measure steps. First, the working electrode was rotated at the desired rate and a cleaning step was performed by cycling 5 times between -0.25 V to 1.25 V versus Ag/AgCl at 200 mV s−1 . Measurements were then performed (while still rotating) by scanning the cell potential between 0.225 and 0.725 V versus Ag/AgCl. The experimental scan rate was chosen empirically from the range of 5–20 mV s−1 to obtain the lowest amount of hysteresis between the forward and reverse sweeps. The voltage range was selected as the minimum needed to obtain steady-state anodic and cathodic limiting currents. This sequence of cleaning and measurement steps was then repeated at each rotation rate of interest. Rotation rates were varied in the order 100, 400, 900, 1600, 2500, 225, 625, 1225, 2025 and 100 rpm, and the experimental data were taken to be valid only if the first and last cycles were found to overlay. After experimentation with Fe3+/2+ electrolyte was complete, background measurements were completed by repeating the entire glassware cleaning protocol and replacing the electrolyte with fresh 0.5 M HCl solution. The aforementioned N2 (g) sparge and electrode cleaning protocols were repeated once more, and then background data were collected by running cyclic voltammetry from 0.225 to 0.725 V versus Ag/AgCl without electrode rotation at the same scan rate as the measurements taken in the presence of RFB active species. These background data were then subtracted from the experimental data prior to analysis. We performed Koutecky-Levich (KL) analysis to find the diffusivity, D, and to extract transport-free current versus overpotential data. This analysis relies on the KL Equation 1, which expresses the relationship between current density and rotation rate in an RDE experiment: 1 1 1 = + ω −0.5 j jk 0.620nF CD2/3 ν −1/6

(1)

where j is the measured current density, jk is kinetic current density, n is number of electrons 8


Page 9 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


transferred (1 in the case of Fe3+/2+ ), F is Faraday’s constant (96485 C/mol), C is bulk reactant concentration, D is diffusivity, ν is kinematic viscosity (taken as 0.01 cm2 s−1 for aqueous solutions at room temperature), and ω is rotation rate in radians/second. Linearized plots of j −1 vs. ω −0.5 , also called KL plots, can be used to extract D from the slope and jk from the y-intercept. We used a mathematical model to fit jk as a function of overpotential, η, using a nonlinear least-squares approach to find the exchange current density, j0 , and the symmetry factors, αox and αred , according to Equation 2.

jk = j0 exp

αox nF η RT

− exp

−αred nF η RT

(2)

where R is the universal gas constant, T is temperature, and η = E − Eeq where Eeq is the empirically measured equilibrium potential of the system. This equation is derived from the ButlerVolmer description of a 1-electron transfer process in the specific case where the oxidized and reduced forms of the redox couple are equal in concentration, which allows for the use of a single j0 term to describe both half-reactions. For a highly uniform electrode surface and a reaction mechanism that is first-order in the electroactive species, j0 can be expressed as a function of reactant concentration, C, and the heterogeneous electron transfer rate constant, k 0 , according to Equation 3. j0 = nF Ck 0

(3)

Results and Discussion A qualitative picture of the kinetics of Fe3+/2+ is apparent from voltammograms under quiescent conditions. Figure 3 depicts representative cyclic voltammetry (CV) data collected at 0 rpm and a scan rate of 200 mV s−1 for Pt and Au electrodes. Oxidative and reductive waves were wellresolved, and the peak-to-peak separation values were 0.08 and 0.14 V for Pt and Au electrodes, respectively, at this scan rate. The electron-transfer can therefore be taken as quasi-reversible in 9



both cases, with Pt exhibiting somewhat faster kinetics. 72 More extensive scan rate dependence data for quiescent CV experiments are included in the Supporting Information, and they show the expected linear increase in peak current with the square root of scan rate for a diffusional redox process.

Figure 3: Representative current density versus applied potential data for (a) polycrystalline Pt and (b) polycrystalline Au in 5mM FeCl2 and 5 mM FeCl3 in 0.5 M HCl(aq) collected at 0 rpm and at scan rate of 200 mV/s. Figure 4 compiles background-subtracted RDE current density versus applied potential (j–E) data for each electrode type over a range of rotation rates from 100 to 2500 rpm. Each of these datasets exhibits a consistent open-circuit potential of 0.475 V versus Ag/AgCl (0.684 V versus NHE), which is close to the reported standard reduction potential of 0.7 versus NHE for Fe3+/2+ redox couple in 1 M HCl electrolyte. 72 Oxidative and reductive current densities increased monotonically in magnitude with increasing rotation rate, as expected for progressively diminished transport limitations. Steady-state limiting current densities were also clearly obtained except at the highest rotation rates in the oxidizing direction, which showed slightly sloping j–E behavior. Figure 5 depicts representative KL analysis, comprising plots of j −1 versus ω −0.5 for overpotentials from 0.02 to 0.14 V in the positive and negative directions. These data exhibited linear trends at overpotentials exceeding 0.02 V in magnitude. KL data collected at smaller overpotentials exhibited excessive spread, which is consistent with increased uncertainty due to mea10


Page 10 of 31

Page 11 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 4: Representative current density versus potential data for RDE voltammetry between 100 and 2500 rpm using (a) polycrystalline Pt and (b) polycrystalline Au in 5mM FeCl2 and 5mM FeCl3 in 0.5 M HCl(aq) at a scan rate of 5 mV/s. surement noise and background subtraction artifacts when the observed current densities were low. The slopes of the best fit lines through the data also converged at overpotentials exceeding 0.06 V, which is consistent with the transition from primarily kinetic to primarily diffusion limitation. Thus, the slopes of the data collected at η ≥ 100 mV were averaged to calculate diffusivity. Figure 6 shows transport-free polarization data for Pt and Au electrodes extracted from linear fits to the KL equation. These data are plotted as natural log of kinetic current density versus overpotential for both the oxidative and reductive half reactions. They qualitatively agree with a Butler-Volmer description of electrode kinetics, as the data show an increasingly linear trend at large overpotentials and deflect to smaller current densities at smaller overpotentials; this deflection is attributable to competition between forward and reverse reactions near equilibrium. Interestingly, application of a conventional Butler-Volmer model where the symmetry factors (αox and αred ) were constrained to sum to 1 for the oxidative and reductive half reactions resulted in rather poor fits, as shown in the Supporting Information. Thus, the fits depicted as black lines in Figure 6 correspond to least squares regression where the α values were allowed to vary independently between 0 and 1 for each half reaction. These fits also included only overpotentials 11



Figure 5: Koutecky-Levich analysis depicting inverse current versus inverse square root of rotation rate data for (a) iron oxidation at Pt, (b) iron reduction at Pt, (c) iron oxidation at Au and (d) iron reduction at Au for overpotentials from 20 to 140 mV in 5mM FeCl2 and 5 mM FeCl3 in 0.5 M HCl.

12


Page 12 of 31

Page 13 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


exceeding 40 mV in magnitude, due to the aforementioned uncertainty implicit in KL analysis of low-overpotential data.

Figure 6: Transport-free polarization data for Fe3+/2+ redox chemistry at (a) Pt and (b) Au electrodes in 5mM FeCl2 and 5 mM FeCl3 in 0.5 M HCl. Fit lines and the noted j0 and α were found using nonlinear least-squares regression using a Butler-Volmer kinetic model.

Table 1 summarizes the extracted transport and kinetics properties of the Fe3+/2+ redox couple on Pt and Au electrodes along with the associated experimental uncertainties. Error bounds were taken at 1 standard deviation (σ) from the mean of 5 replicates. We found that the diffusivity Table 1: Transport and kinetics properties of iron oxidation and reduction at Pt and Au electrodes. Uncertainty values (1σ) are reported in parentheses on the basis of n=5 sets of experiments electrode Pt (n=5) Au (n=5)

Dox (cm2 s−1 ) 4 x 10−6 (±0.1 x 10−6 ) 4 x 10−6 (±0.5 x 10−6 )

Dred (cm2 s−1 ) 3 x 10−6 (±0.2 x 10−6 ) 3 x 10−6 (±0.2 x 10−6 )

j0 (mA cm−2 ) 3.7 (±0.5) 1.3 (±0.2)

αox 0.33 (±0.08) 0.25 (±0.05)

αred 0.35 (±0.09) 0.45 (±0.05)

values for Pt and Au were between 3–4 x 10−6 cm2 s−1 , and they were indistinguishable within a single standard deviation. However, exchange current densities were clearly different: 3.7 ± 0.5 mA cm−2 for polycrystalline Pt and 1.3 ± 0.2 mA cm−2 for polycrystalline Au. Thus, the Fe3+/2+ redox reaction was indeed faster at Pt than Au, but only by a factor of 3. Finally, when symmetry

13



factors were allowed to vary arbitrarily, they were found to be α < 0.5 in all cases. The αox for Pt and Au was between 0.25 and 0.33, while αred was between 0.35 and 0.45. RDE voltammetry is attractive for kinetics measurements of RFBs since it is already well established as an electroanalytical method. In fact, this method has been widely used in recent studies of RFB active materials. 37–39,73–75 RDE voltammetry also has the advantage of being commonly used as an analytical protocol in electrocatalysis. 68–71,76,77 As a result, instrumentation and expertise is readily translatable between these two fields. Nevertheless, it remains unclear whether these types of analytical methods can be convincingly used to predict or explain the behavior of functional RFB devices. The present study represents an initial effort in our laboratory to address this key challenge. Our experimental approach comprised an effort to balance between the goals of mimicking functional flow battery behavior while also maintaining suitable analytical precision. Aqueous Fe chloride salts and HCl supporting electrolyte were chosen because they form the basis of the Fe positive electrolyte in the Fe/Cr RFB. The redox couple was added at relatively low concentration—10 mM total Fe—while the supporting electrolyte was maintained at 0.5 M. This diverges from electrolyte conditions in a functional flow battery to avoid the confounding effects attributable to uncompensated resistance that arise from high current densities in these systems. We also chose to study equimolar mixtures of Fe2+ and Fe3+ , which simplifies the experimental procedure by enabling analysis of oxidative and reductive half-reactions in a single series of RDE voltammograms. Moreover, a mixture of oxidized and reduced forms of the redox couple resembles the conditions of a working RFB cell, since these batteries are rarely cycled to 0 or 100 % state of charge once in operation. Although Pt and Au are not widely used as electrodes in functional RFBs (due to their high cost), they were chosen for this study because there exist well-established preparation and cleaning protocols for these materials. RDE voltammetry also approximates the hydrodynamic conditions in an RFB. The diffusion boundary at a rotating electrode is in fact more homogeneous than in a flow cell, thereby permitting straightforward use of KL analysis to extract kinetic and transport parameters. 14


Page 14 of 31

Page 15 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Although the cell and electrode preparation procedures we used are onerous, we found that every step was necessary to successfully differentiate the catalytic activities of Pt and Au toward Fe3+/2+ . To illustrate, Figure 7 shows the results of RDE voltammetry experiments at a Pt electrode where only one step—the electrochemical cleaning step implemented between collecting voltammograms at different rotation rates—was omitted. In this case, we observed a clear progressive decrease in the current response at a given applied potential over time, which is particularly evident when the rotation rates were varied in the order shown in Figure 7. We attribute the decreased current to progressive electrode fouling, which was mitigated by electrochemically cleaning the electrode between each RDE measurement. It is relevant to note that, while cleaning steps like those employed here are routinely implemented in the course of electroanalysis, it would not be practical to implement an analogous cleaning protocol in a functional RFB. Thus, it may not be reasonable to expect that electrode pre-treatments will give rise to optimal kinetics over extended battery operation.

Figure 7: (a) RDE current density versus potential data using a Pt electrode in 5 mM FeCl2 and 5 mM FeCl3 in 0.5 M HCl obtained by omitting cleaning steps between different rotation rates. (b) Current density versus rotation rate at 375 mV versus Ag/AgCl with and without cleaning steps. The numbered dotted lines represent the order in which RDE measurements were performed.

In studies of electrochemical kinetics, it is common to report a heterogeneous electrontransfer rate constant, k 0 , as the primary descriptor. This parameter is directly analogous to the 15



characteristic rate constant k in solution-phase chemical kinetics. However, unlike homogeneous reactions in which the reacting species are highly uniform in composition and structure, polycrystalline electrode surfaces exhibit considerable surface site heterogeneity. Thus, for reactions in which the electrode surface acts as a catalyst, each site likely exhibits a different characteristic k 0 , and these may vary widely. This helps explain why consistent electrode surface preparation is so important to obtain reproducible kinetics data. We further conjecture that consistent application of two different surface preparation protocols may give rise to highly precise and reproducible kinetics measurements that nonetheless diverge from one another. Thus, we recommend against the use of k 0 to describe RFB electron-transfer kinetics unless the electrode surface is highly uniform in chemical composition and structure, such as when using single crystals. We have instead reported j0 values, which are already commonly used in electrocatalysis research and clearly depend on extensive properties of the electrode (surface area, density of active sites) and the electrolyte (reactant concentration). It would also be reasonable to report apparent k 0 values along with detailed information about electrolyte composition and electrode surface preparation. To further assess the utility of our analytical approach for predicting the practical performance of functional flow batteries, we have carried out RDE analysis of polycrystalline Pt in a more realistic Fe RFB positive electrolyte containing 1 M total Fe (0.5 M each of FeCl2 and FeCl3 ) in 2 M HCl(aq). All analytical procedures were followed exactly as described in the Experimental section, except for the change in electrolyte composition and the compensation of solution resistance, which was measured to be 4 ohms via impedance spectroscopy. Figure 8 shows representative RDE and transport-free polarization data. We observed a negative shift in the equilibrium potential to ∼0.425 V versus Ag/AgCl in the concentrated Fe RFB electrolyte, which is consistent with increased free chloride concentration. Thus, the thermodynamics of the reaction depend on the extent to which Cl – dissociates from the Fe salts as the electrolyte concentration is increased. We also found j0 = 60 mA cm−2 , which is ∼6-fold lower than the value of 370 (±50) mA cm−2 that would be inferred from the simple extrapolation of the low-concentration data based on the linear relationship between j0 and reac16


Page 16 of 31

Page 17 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 8: (a) RDE current density versus potential data for a Fe3+/2+ RFB electrolyte containing 0.5 M FeCl3 and 0.5 M FeCl2 in 2 M HCl(aq). Data were collected at a scan rate of 5 mV s−1 and series resistance was compensated at 4 Ω. The inset shows the same data without compensating for series resistance. (b) The corresponding transport free polarization data, where the equilibrium potential Eeq was taken to be 425 mV versus Ag/AgCl. tant concentration for a process that is first-order in Fe. We also found the analysis of these data to be sensitive to the specifics of series resistance compensation; for several datasets, compensation at the experimentally measured value of 4 Ω resulted in un-physical results in KL analysis, while omitting the resistance compensation entirely resulted in j0 values that were lower by approximately a factor of 2. Thus, we ascribe somewhat larger uncertainties to these high-concentration data. These results provide considerable impetus for further work aimed at understanding the differences in thermodynamics and kinetics of RFB redox chemistry under analytical (low concentration) and practical (high concentration) conditions. Although the Fe3+/2+ redox couple involves the transfer of only one electron, our results are broadly consistent with a multi-step reaction mechanism where the electrode surface plays a catalytic role. Two readily apparent indications are that (a) Pt and Au electrodes clearly showed different electron-transfer kinetics, and (b) even at Pt, the reaction kinetics were considerably slower than those routinely observed from outer-sphere, 1-electron transfer reagents like ferrocene. 78 Moreover, acceptable fits to transport-free polarization data did not obey a classi17



cal BV model where the symmetry factors for the forward and reverse reactions sum to one. This characteristic of the BV model is based on the assumption of a mechanism where a single electrontransfer step governs the rate of the forward and reverse reactions, whereas α values that diverge from the expected relationship are consistent with multi-step redox chemistry. Prior work on aqueous Fe3+/2+ redox chemistry found that multi-step reaction pathways involving bridging anions accompany electron transfer in homogeneous solution. 79–84 Surfaceadsorbed anions have also been strongly implicated as catalysts in the electrochemistry of Fe aquo complexes. 50–52,59 Thus, the HCl supporting electrolyte used in Fe/Cr RFBs is desirable for its low cost and ability to solubilize Fe and Cr salts, and it may also play an advantageous catalytic role. To address this possibility, we carried out additional RDE measurements of Fe3+/2+ redox chemistry in HClO4 and H2 SO4 supporting electrolytes (see Supporting Information). The apparent reaction kinetics increased in the order H2 SO4 ≈ HClO4 < HCl, but the differences in overpotential performance for polycrystalline Pt in all three supporting electrolytes were very small. These results suggest that chloride and sulfate adsorbates are comparably effective at catalyzing the redox chemistry of Fe aquo complexes. The observation of similar reaction rates in HClO4 is surprising since ClO4 – is a non-adsorbing anion. However, Weber previously observed that extremely small impurity concentrations of Cl – were sufficient to accelerate the kinetics of aquo Fe3+/2+ redox chemistry nearly as much as in more concentrated chloride electrolytes. 52 Our observations are consistent with this result, as we did not make the necessary efforts to entirely exclude chloride impurities. Improved insights into RFB interfacial electron transfer dynamics will depend critically on developing consistent preparation methods not just for model electrode surfaces like Pt and Au, but also for technologically relevant RFB electrode materials like graphitic carbon in various forms. However, carbon electrodes pose an additional challenge: they are less straightforward than noble metals to clean due to their surface redox reactivity. To illustrate this point, Figure 9 depicts cyclic voltammograms of 10 mM Fe3+/2+ in 0.5 M HCl(aq) using glassy carbon working electrodes whose surfaces were polished and then subjected to three different surface cleaning protocols. The first involved submersing the electrode in isopropanol that had been pre-purified with activated 18


Page 18 of 31

Page 19 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 9: Current density versus potential data for glassy carbon working electrodes in 5mM FeCl2 and 5 mM FeCl3 in 0.5 M HCl at a scan rate of 200 mV/s. Prior to measurement, the electrodes were polished and then subjected to different surface treatments: (a) solvent cleaning via submersion for 10 minutes in isopropanol that had been purified with activated carbon; (b) isopropanol cleaning followed by electrochemical cycling between -0.25 and 1.7 V vs Ag/AgCl for ∼3 minutes in 0.5 M H2 SO4 ; and (c) isopropanol cleaning followed by submersion in 30 wt% H2 O2 (aq) for 10 minutes. carbon. This treatment has been reported as effective for removing oxygen-containing surface contaminants from glassy carbon surfaces that persist even after polishing. 85 The second treatment involved the same isopropanol cleaning step, followed by extended electrochemical cycling in aqueous H2 SO4 solution between -0.25 and 1.7 V vs Ag/AgCl. This is a common method to clean and oxidatively activate carbon electrodes. 86 The third treatment again involved isopropanol cleaning followed by submersion of the electrode surface in 30 wt% aqueous H2 O2 for 10 minutes, which is a straightforward technique for oxidizing carbon surfaces. 87 These results are remarkable because the isopropanol and H2 O2 -treated surfaces each exhibited extremely sluggish kinetics for 1-electron Fe3+/2+ redox chemistry, as evidenced by their large ∆Ep values. Moreover, while the electrochemical and H2 O2 treatments both oxidize carbon surfaces, the former resulted in greatly enhanced electron-transfer kinetics (qualitatively comparable to Au) while the latter resulted in slightly decreased kinetics. Additional work is currently underway to probe the relationship between carbon electrode preparation conditions, surface chemistry, and electron-transfer kinetics.

19



Conclusions The primary aim of this study was to deploy reproducible methods to characterize RFB electron transfer kinetics under conditions that are analytically precise and technologically relevant. Our approach was based on RDE techniques, for which instrumentation and expertise are widely available. We obtained reproducible results for Fe3+/2+ in HCl(aq), but only if rigorous steps were performed to maintain scrupulously clean surfaces. Our results also clearly indicated that the electrode surface plays a role, which is consistent with prior electroanalytical work showing that adsorbed anions catalyze Fe redox chemistry. Further work is warranted to understand changes in kinetics that occur as RFB electrolyte concentration is increased, and also to adapt these methods for the complex surface chemistry of carbon electrodes. To this end, ongoing work in our lab is focused on the development of electroanalytical methods that more closely resemble functional RFB architectures. Overall, the results from this study highlight the importance of electrochemical catalysis in the design of RFBs and therefore provide a basis for continued convergence between research in batteries and fuel cell energy storage.

Acknowledgements We gratefully acknowledge the Swanson School of Engineering at the University of Pittsburgh for support of this work via startup funds for the McKone Laboratory. We also acknowledge Rituja Patil, Eli Bostian, Yifan Deng, Aayush Mantri, and Emily Siegel for their editorial feedback during the preparation of this manuscript.

20


Page 20 of 31

Page 21 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Supporting Information Available Detailed description of the simulations of RFB overpotential performance. Determination of cell purge time. Evaluation of electrode surface roughness. Tabulated Fe3+/2+ kinetics data from previous literature reports. CV scan rate dependence data. Results of Butler-Volmer fits with constrained α values. RDE analysis in different supporting electrolytes.

This material is available

free of charge via the Internet at http://pubs.acs.org/.

References (1) Johansson, T. B.; Kelly, H.; Reddy, A. K.; Willliams, R. Renewable Fuels and Electricity for a Growing World Economy: Defining and Achieving the potential. Energy Studies Review 1992, 4, 201–222. (2) Dincer, I. Renewable Energy and Sustainable Development: A Crucial Review. Renewable and Sustainable Energy Reviews 2000, 4, 157–175. (3) Luo, X.; Wang, J.; Dooner, M.; Clarke, J. Overview of Current Development in Electrical Energy Storage Technologies and the Application Potential in Power System Operation. Appl. Energy 2015, 137, 511–536. (4) Hill, C. A.; Such, M. C.; Chen, D.; Gonzalez, J.; Grady, W. M. K. Battery Energy Storage for Enabling Integration of Distributed Solar Power Generation. IEEE Transactions on Smart Grid 2012, 3, 850–857. (5) Moslehi, K.; Kumar, R. A Reliability Perspective of the Smart Grid. IEEE Transactions on Smart Grid 2010, 1, 57–64. (6) Tarroja, B.; Mueller, F.; Eichman, J. D.; Brouwer, J.; Samuelsen, S. Spatial and Temporal Analysis of Electric Wind Generation Intermittency and Dynamics. Renewable Energy 2011, 36, 3424 – 3432. 21



(7) Chu, S.; Majumdar, A. Opportunities and Challenges for a Sustainable Energy Future. Nature 2012, 488, 294–303. (8) Dell, R. M.; Rand, D. A. J. Energy Storage - A Key Technology for Global Energy Sustainability. J. Power Sources 2001, 100, 2 – 17. (9) Dunn, B.; Kamath, H.; Tarascon, J. Electrical Energy Storage for the Grid: A Battery of Choices. Science 2011, 334, 928–935. (10) Soloveichik, G. L. Battery Technologies for Large-Scale Stationary Energy Storage. Annu. Rev. Chem. Biomol. Engg. 2011, 2, 503–527. (11) Jens, N.; Nataliya, R.; Tatjana, H.; Peter, F. The Chemistry of Redox Flow Batteries. Angew. Chem., Int. Ed. 2015, 54, 9776–9809. (12) Tokuda, N.; Kumamoto, T.; Shigematsu, T.; Deguchi, H.; Ito, T.; Yoshikawa, N.; Hara, T. Development of a Redox Flow Battery System. Sei Tech Rev 1998, 88–94. (13) Leung, P.; Li, X.; de Leon, C.; Berlouis, L.; Low, C. T. J.; Walsh, F. C. Progress in Redox Flow Batteries, Remaining Challenges and their Applications in Energy Storage. RSC Adv. 2012, 2, 10125–10156. (14) Nguyen, T.; Savinell, R. F. Flow Batteries. Electrochem. Soc. 2010, 19, 54–56. (15) Darling, R. M.; Gallagher, K. G.; Kowalski, J. A.; Ha, S.; Brushett, F. R. Pathways to LowCost Electrochemical Energy Storage: A Comparison of Aqueous and Nonaqueous Flow Batteries. Energy Environ. Sci. 2014, 7, 3459–3477. (16) Bartolozzi, M. Development of Redox Flow Batteries. A Historical Bibliography. J. Power Sources 1989, 27, 219 – 234. (17) Pan, F.; Wang, Q. Redox Species of Redox Flow Batteries: A Review. Molecules 2015, 20, 20499 – 20517. 22


Page 22 of 31

Page 23 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(18) Wei, W.; Qingtao, L.; Bin, L.; Xiaoliang, W.; Liyu, L.; Zhenguo, Y. Recent Progress in Redox Flow Battery Research and Development. Adv. Functi. Mater. 2012, 23, 970–986. (19) Yang, Z.; Zhang, J.; Kintner-Meyer, M. C. W.; Lu, X.; Choi, D.; Lemmon, J. P.; Liu, J. Electrochemical Energy Storage for Green Grid. Chem. Rev. 2011, 111, 3577–3613. (20) Weber, A. Z.; Mench, M. M.; Meyers, J. P.; Ross, P. N.; Gostick, J. T.; Liu, Q. Redox Flow Batteries: A Review. J. Appl. Electrochem. 2011, 41, 1137–1164. (21) Skyllas-Kazacos, M.; Chakrabarti, M. H.; Hajimolana, S. A.; Mjalli, F. S.; Saleem, M. Progress in Flow Battery Research and Development. J. Electrochem. Soc. 2011, 158, R55– R79. (22) de León, C. P.; Fr´ıas-Ferrer, A.; González-Garc´ıa, J.; Szánto, D. A.; Walsh, F. C. Redox Flow Cells for Energy Conversion. J. Power Sources 2006, 160, 716 – 732. (23) Sauer, D. U.; Wenzl, H. Comparison of Different Approaches for Lifetime Prediction of Electrochemical Systems Using Lead-Acid Batteries as Example. J. Power Sources 2008, 176, 534–546. (24) Soloveichik, G. L. Flow Batteries: Current Status and Trends. Chem. Rev. 2015, 115, 11533– 11558. (25) Hagedorn, N.; Thaller, L. Redox Storage Systems for Solar Applications. NASA TM 81464, National Aeronautics and Space Administration, U S Dept. of Energy 1980, (26) Gahn, R. F.; Hagedorn, N. H. Single Cell Performance Studies on the Fe / Cr Redox Energy Storage System Using Mixed Reactant Solutions at Elevated Temperature Conservation and Renewable Energy Division of Energy Storage Systems. NASA TM-83385, National Aeronautics and Space Administration, U S Dept. of Energy 1983, (27) Thaller, L. H. Electrically Rechargeable Redox Flow Cell. United States Patent 3,996,064 1976, 23



(28) Kaneko, H.; Nozaki, K.; Wada, Y.; Aoki, T.; Negishi, A.; Kamimoto, M. Vanadium Redox Reactions and Carbon Electrodes for Vanadium Redox Flow Battery. Electrochim. Acta 1991, 36, 1191–1196. (29) Maria, S.; George, K.; Grace, P.; Hugh, V. Recent Advances with UNSW Vanadium Based Redox Flow Batteries. Int. J. Energy Res. 2010, 34, 182–189. (30) Skyllas-Kazacos, M. Novel Vanadium Chloride/Polyhalide Redox Flow Battery. J. Power Sources 2003, 124, 299–302. (31) Rychcik, M.; Skyllas-Kazacos, M. Characteristics of a New All-Vanadium Redox Flow Battery. J. Power Sources 1988, 22, 59–67. (32) Rychcik, M.; Skyllas-Kazacos, M. Evaluation of Electrode Materials for Vanadium Redox Cell. J. Power Sources 1987, 19, 45–54. (33) Sun, B.; Skyllas-Kazacos, M. Modification of Graphite Electrode Materials for Vanadium Redox Flow Battery Application I. Thermal Treatment. Electrochim Acta 1992, 37, 1253– 1260. (34) Sun, B.; Skyllas-Kazacos, M. Chemical Modification of Graphite Electrode Materials for Vanadium Redox Flow Battery Application Part II. Acid Treatments. Electrochim Acta 1992, 37, 2459–2465. (35) Aaron, D. S.; Liu, Q.; Tang, Z.; Grim, G. M.; Papandrew, A. B.; Turhan, A.; Zawodzinski, T. A.; Mench, M. M. Dramatic Performance Gains in Vanadium Redox Flow Batteries through Modified Cell Architecture. J. Power Sources 2012, 206, 450–453. (36) Houser, J.; Clement, J.; Pezeshki, A.; Mench, M. M. Influence of Architecture and Material Properties on Vanadium Redox Flow Battery Performance. J. Power Sources 2016, 302, 369– 377.

24


Page 24 of 31

Page 25 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(37) Lin, K.; Gómez-Bombarelli, R.; Beh, E. S.; Tong, L.; Chen, Q.; Valle, A.; Aspuru-Guzik, A.; Aziz, M. J.; Gordon, R. G. A Redox-Flow Battery with an Alloxazine-Based Organic Electrolyte. Nat. Energy 2016, 1, 1–8. (38) Huskinson, B.; Marshak, M. P.; Suh, C.; Er, S.; Gerhardt, M. R.; Galvin, C. J.; Chen, X.; Aspuru-Guzik, A.; Gordon, R. G.; Aziz, M. J. A Metal-Free Organic - Inorganic Aqueous Flow Battery. Nature 2014, 505, 195–198. (39) Tianbiao, L.; Xiaoliang, W.; Zimin, N.; Vincent, S.; Wei, W. A Total Organic Aqueous Redox Flow Battery Employing a Low Cost and Sustainable Methyl Viologen Anolyte and 4HOTEMPO Catholyte. Adv. Energy Mater. 2015, 6, 1–8. (40) Hu, B.; DeBruler, C.; Rhodes, Z.; Liu, T. L. Long-Cycling Aqueous Organic Redox Flow Battery (AORFB) toward Sustainable and Safe Energy Storage. J. Am. Chem. Soc 2017, 139, 1207–1214. (41) Beh, E. S.; De Porcellinis, D.; Gracia, R. L.; Xia, K. T.; Gordon, R. G.; Aziz, M. J. A Neutral pH Aqueous OrganicOrganometallic Redox Flow Battery with Extremely High Capacity Retention. ACS Energy Letters 2017, 2, 639–644. (42) Lin, K.; Chen, Q.; Gerhardt, M. R.; Tong, L.; Kim, S. B.; Eisenach, L.; Valle, A. W.; Hardee, D.; Gordon, R. G.; Aziz, M. J.; Marshak, M. P. Alkaline Quinone Flow Battery. Science 2015, 349, 1529–1532. (43) Gong, K.; Fang, Q.; Gu, S.; Li, S. F. Y.; Yan, Y. Nonaqueous Redox-Flow Batteries: Organic Solvents, Supporting Electrolytes, and Redox Pairs. Energy Environ. Sci. 2015, 8, 3515– 3530. (44) Milshtein, J. D.; Barton, J. L.; Carney, T. J.; Kowalski, J. A.; Darling, R. M.; Brushett, F. R. Towards Low Resistance Nonaqueous Redox Flow Batteries. J. Electrochem. Soc. 2017, 164, A2487–A2499.

25



(45) Suttil, J. A.; Kucharyson, J. F.; Escalante-Garcia, I. L.; Cabrera, P. J.; James, B. R.; Savinell, R. F.; Sanford, M. S.; Thompson, L. T. Metal Acetylacetonate Complexes for High Energy Density Non-Aqueous Redox Flow Batteries. J. Mater. Chem. A 2015, 3, 7929–7938. (46) Sevov, C. S.; Brooner, R. E. M.; Chénard, E.; Assary, R. S.; Moore, J. S.; Rodr´ıguezLópez, J.; Sanford, M. S. Evolutionary Design of Low Molecular Weight Organic Anolyte Materials for Applications in Nonaqueous Redox Flow Batteries. J. Am. Chem. Soc 2015, 137, 14465–14472. (47) Xiaoliang, W.; Wu, X.; Jinhua, H.; Lu, Z.; Eric, W.; Chad, L.; Vijayakumar, M.; Henderson, W. A.; Tianbiao, L.; Lelia, C.; Bin, L.; Vincent, S.; Wei, W. Radical Compatibility with Nonaqueous Electrolytes and Its Impact on an All Organic Redox Flow Battery. Angewandte Chemie International Edition 2015, 54, 8684–8687. (48) Jordan, J.; Javick, R. A. Electrode Kinetics by Hydrodynamic Voltammetry Study of FerrousFerric, Ferrocyanide-Ferricyanide and Iodide-Iodine Systems. Electrochim. Acta 1962, 6, 23– 33. (49) Jahn, D.; Vielstich, W. Rates of Electrode Processes by the Rotating Disk Method. J. Electrochem. Soc. 1962, 109, 849–852. (50) Angell, D. H.; Dickinson, T. The Kinetics of the Ferrous/Ferric and Ferro/Ferricyanide Reactions at Platinum and Gold Electrodes: Part I. Kinetics at Bare-Metal Surfaces. J. Electroanal. Chem. Interfacial Electrochem. 1972, 35, 55–72. (51) Junsuke, S. Hydrodynamic Voltammetry with the Convection Electrode. IV. The Measurements of the Kinetic Parameters of the Electrode Reaction. Bull. Chem. Soc. Jpn. 1970, 43, 755–758. (52) Weber, J.; Samec, Z.; Mareˇcek, V. The Effect of Anion Adsorption on the Kinetics of the Fe(3+)/Fe(2+) Reaction on Pt and Au Electrodes in HClO4. J. Electroanal. Chem. Interfacial Electrochem. 1978, 89, 271–288. 26


Page 26 of 31

Page 27 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(53) Breckenridge, J. H.; Harris, W. E. Effect of Thorium on the Reactions of Iron(II) and Iron(III) at a Platinum Electrode. Can. J. Chem. 1970, 48, 1934–1936. (54) McDermott, C. A. A.; Kneten, K. R.; Mccreery, R. L. Electron Transfer Kinetics of Aquated Fe(+3/+2), Eu(+3/+2), and V(+3/+2) at Carbon Electrodes: Inner Sphere Catalysis by Surface Oxides. J. Electrochem. Soc. 1993, 140, 2593–2599. ˇ ıková, M.; Vydra, F. Voltammetry with Disk Electrodes and its Analytical Application: (55) Stul´ IV. The Voltammetry of Iron(III) at the Glassy Carbon Rotating Disk Electrode in Acid Media. J. Electroanal. Chem. Interfacial Electrochem. 1972, 38, 349–357. (56) Galus, Z.; Adams, R. N. The Investigation of the Kinetics of Moderately Rapid Electrode Reactions Using Rotating Disk Electrodes. J. Phys. Chem 1963, 67, 866–871. (57) Randles, J. E. B. Kinetics of Rapid Electrode Reactions. Discuss. Faraday Soc. 1947, 1, 11– 19. (58) Agarwal, H. P.; Qureshi, S. Faradaic Rectification Studies of the Ferrous-Ferric Redox Couple in the Acid Media. Electrochim. Acta 1974, 19, 607–610. (59) Hung, N. C.; Nagy, Z. Kinetics of the Ferrous/Ferric Electrode Reaction in the Absence of Chloride Catalysis. J. Electrochem. Soc. 1987, 134, 2215–2220. (60) Anson, F. C. The Effect of Surface Oxidation on the Voltammetric Behavior of Platinum Electrodes. The Electroreduction of Iodate. J. Am. Chem. Soc 1959, 81, 1554–1557. (61) Zhong, S.; Skyllas-Kazacos, M. Electrochemical Behaviour of Vanadium(V)/Vanadium(IV) Redox Couple at Graphite Electrodes. J. Power Sources 1992, 39, 1–9. (62) Sum, E.; Rychcik, M.; Skyllas-kazacos, M. Investigation of the V(V)/V(IV) System for use in the Positive Half-Cell of a Redox Battery. J. Power Sources 1985, 16, 85–95.

27



(63) Aaron, D.; Sun, C. N.; Bright, M.; Papandrew, A. B.; Mench, M. M.; Zawodzinski, T. A. In Situ Kinetics Studies in All-Vanadium Redox Flow Batteries. ECS Electrochem. Lett. 2013, 2, A29–A31. (64) Oriji, G.; Katayama, Y.; Miura, T. Investigations on V(IV)/V(V) and V(II)/V(III) Redox Reactions by Various Electrochemical Methods. J. Power Sources 2005, 139, 321–324. (65) Agar, E.; Dennison, C. R.; Knehr, K. W.; Kumbur, E. C. Identification of Performance Limiting Electrode using Asymmetric Cell Configuration in Vanadium Redox Flow Batteries. J. Power Sources 2013, 225, 89–94. (66) Fink, H.; Friedl, J.; Stimming, U. Composition of the Electrode Determines Which HalfCells Rate Constant is Higher in a Vanadium Flow Battery. J. Phys. Chem. C 2016, 120, 15893–15901. (67) Aaron, D.; Tang, Z.; Papandrew, A. B.; Zawodzinski, T. A. Polarization Curve Analysis of All-Vanadium Redox Flow Batteries. J. Appl. Electrochem. 2011, 41, 1175–1182. (68) Garsany, Y.; Baturina, O. A.; Swider-Lyons, K. E.; Kocha, S. S. Experimental Methods for Quantifying the Activity of Platinum Electrocatalysts for the Oxygen Reduction Reaction. Anal. Chem 2010, 82, 6321–6328. (69) Gasteiger, H. A.; Kocha, S. S.; Sompalli, B.; Wagner, F. T. Activity Benchmarks and Requirements for Pt, Pt-alloy, and non-Pt Oxygen Reduction Catalysts for PEMFCs. Appl. Catal., B 2005, 56, 9–35. (70) McCrory, C. C. L.; Jung, S.; Peters, J. C.; Jaramillo, T. F. Benchmarking Heterogeneous Electrocatalysts for the Oxygen Evolution Reaction. J. Am. Chem. Soc 2013, 135, 16977– 16987. (71) McCrory, C. C. L.; Jung, S.; Ferrer, I. M.; Chatman, S. M.; Peters, J. C.; Jaramillo, T. F.

28


Page 28 of 31

Page 29 of 31 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Benchmarking Hydrogen Evolving Reaction and Oxygen Evolving Reaction Electrocatalysts for Solar Water Splitting Devices. J. Am. Chem. Soc 2015, 137, 4347–4357. (72) J Bard, A.; R. Faulkner, L. Electrochemical Methods : Fundamentals and Applications, 2nd ed.; John Wiley and Sons, Inc., 1980; Vol. 18. (73) Zhengjin, Y.; Liuchuan, T.; Tabor Daniel, P.; Beh, E. S.; Marc-Antoni, G.; Diana, P.; Alán, A.G.; Gordon, R. G.; Aziz, M. J. Alkaline Benzoquinone Aqueous Flow Battery for Large-Scale Storage of Electrical Energy. Adv. Energy Mater. 8, 1–9. (74) Finkelstein, D. A.; Mota, N. D.; Cohen, J. L.; Abruña, H. D. Rotating Disk Electrode (RDE) Investigation of BH4- and BH3OH- Electro-oxidation at Pt and Au: Implications for BH4Fuel Cells. J. Phys. Chem. C 2009, 113, 19700–19712. (75) Wakabayashi, R. H.; Paik, H.; Murphy, M. J.; Schlom, D. G.; Brützam, M.; Uecker, R.; van Dover, R. B.; DiSalvo, F. J.; Abruña, H. D. Rotating Disk Electrode Voltammetry of Thin Films of Novel Oxide Materials. J. Electrochem. Soc. 2017, 164, H1154–H1160. (76) Janoschka, T.; Martin, N.; Martin, U.; Friebe, C.; Morgenstern, S.; Hiller, H.; Hager, M. D.; Schubert, U. S. An Aqueous, Polymer-Based Redox-Flow Battery using Non-Corrosive, Safe, and Low-Cost Materials. Nature 2015, 527, 78–81. (77) Ding, Y.; Zhao, Y.; Li, Y.; Goodenough, J. B.; Yu, G. A High-Performance All-MetalloceneBased, Non-Aqueous Redox Flow Battery. Energy Environ. Sci. 2017, 10, 491–497. (78) Pournaghi-Azar, M. H.; Ojani, R. Electrode Kinetic Parameters of the Ferrocene Oxidation at Platinum, Gold and Glassy Carbon Electrodes in Chloroform. Electrochim. Acta 1994, 39, 953–955. (79) Rosso, K. M.; Smith, D. M. A.; Dupuis, M. Aspects of Aqueous Iron and Manganese (II/III) Self-Exchange Electron Transfer Reactions. J. Phys. Chem. A 2004, 108, 5242–5248.

29



(80) Brunschwig, B. S.; Logan, J.; Newton, M. D.; Sutin, N. A Semiclassical Treatment of Electron-Exchange Reactions. Application to the Hexaaquoiron(II)-Hexaaquoiron(III) System. J. Am. Chem. Soc 1980, 102, 5798–5809. (81) Silverman, J.; Dodson, R. W. The Exchange Reaction between the Two Oxidation States of Iron in Acid Solution. J. Phys. Chem 1952, 56, 846–852. (82) Jolley, W. H.; Stranks, D. R.; Swaddle, T. W. Pressure Effect on the Kinetics of the Hexaaquairon(II/III) Self-Exchange Reaction in Aqueous Perchloric Acid. Inorg. Chem. 1990, 29, 1948–1951. (83) Logan, J.; Newton, M. D. Ab Initio Study of Electronic Coupling in the Aqueous Fe2+ - Fe3+ Electron Exchange Process. J. Chem. Phys. 1983, 78, 4086–4091. (84) Tembe, B. L.; Friedman, H. L.; Newton, M. D. The Theory of the Fe2+ - Fe3+ Electron Exchange in Water. J. Chem. Phys. 1982, 76, 1490–1507. (85) Ranganathan, S.; Kuo, T. C.; McCreery, R. L. Facile Preparation of Active Glassy Carbon Electrodes with Activated Carbon and Organic Solvents. Anal. Chem 1999, 71, 3574–3580. (86) Kinoshita, K. Carbon : electrochemical and physicochemical properties; Wiley: New York, 1988; p 316. ´ ao, J. J. M. Modification of the (87) Figueiredo, J. L.; Pereira, M. F. R.; Freitas, M. M. A.; Orf˜ Surface Chemistry of Activated Carbons. Carbon 1999, 37, 1379–1389.

30


Page 30 of 31

Page 31 of 31

Graphical TOC Entry 2

mA/cm2

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1 0

-1

2+ Fe(aq)

3+ Fe(aq)

-2 0.3

0.4

0.5

0.6

V vs. Ag/AgCl

31


0.7

Flow Battery Electroanalysis: Hydrodynamic Voltammetry of Aqueous

Recommend Documents