Multielectron Cycling of a Low-Potential Anolyte in Alkali Metal Electrolytes for Nonaqueous Redox Flow Batteries Koen H. Hendriks,§,†,‡ Christo S. Sevov,§,†,‡ Monique E. Cook,†,‡ and Melanie S. Sanford*,†,‡ †
Department of Chemistry, University of Michigan, 930 North University Avenue, Ann Arbor, Michigan 48109, United States Joint Center for Energy Storage Research (JCESR), Argonne, Illinois 60439, United States
‡
S Supporting Information *
ABSTRACT: Recent efforts have led to the design of new anolytes for nonaqueous flow batteries that exhibit reversible redox couples at low potentials. However, these molecules generally cycle through just a single electron-transfer event, which limits the overall energy density of resulting batteries on account of the undesirably high equivalent weight (i.e., ratio of anolyte/supporting electrolyte molecular weight to electrons transferred). In addition, these anolytes generally require expensive alkylammonium salts as supporting electrolytes for stable cycling, which further increases the equivalent weight of the system. The current work describes the multielectron redox cycling of a low-potential anolyte using alkali metal salts as supporting electrolytes. These studies reveal that potassium hexafluorophosphate (KPF6) dramatically lowers the equivalent weight of the anolyte system while supporting flow cell cycling through two redox events at low potentials for 150 cycles with no detectable degradation. edox flow batteries (RFBs) are energy storage devices that can assist with the large-scale integration of renewable energy sources into the electrical grid.1−4 RFBs store energy in solvated redox-active molecules, thus enabling battery capacity to be decoupled from power.5−9 As such, flow batteries are more cost-effective to scale than those with solid-state architectures. The largest components of gridscale RFBs are the electrolyte solutions. To lower the cost of these solutions, techno-economic (TE) models indicate that the storage materials must have high energy densities (>25 Wh/L),10,11 which can be achieved by maximizing the opencircuit voltage (OCV), the concentration of redox-active compounds, and the number of transferrable electrons per molecule. Additionally, the TE models predict that the cost of redox active molecules and supporting electrolytes is generally proportional to their equivalent weights (i.e., ratios of molecular weight to the number of transferable electrons). As such, TE models set an equivalent weight target of 150 g mol−1 e−1 for redox active molecule−supporting electrolyte pairs. To this end, our group has focused on developing redox active organic molecules that undergo reversible redox at potentials that exceed the voltaic stability limits of water.12,13 For example, using a combination of synthetic modification and predictive modeling, we developed the 4-benzoylpyridinium derivative 1+ as a RFB anolyte (eq 1).14 Charge−discharge experiments in MeCN with tetrabutylammonium hexafluor-
R
© XXXX American Chemical Society
ophosphate (TBAPF6) as the supporting electrolyte demonstrated stable one-electron cycling (1+ ⇄ 1) over 200 cycles. While these preliminary results are encouraging, the oneelectron cycling of 1+ in TBAPF6 yields an undesirably high equivalent weight of 312.8 g mol−1 e−1. This is more than double the limit for an economically viable storage material.
We sought to address this challenge by focusing on two variables that contribute to the equivalent weight of 1+: (1) the number of electrons transferred per molecule and (2) the molecular weight of the supporting electrolyte. First, accessing two-electron transfers (1+ ⇄ 1 ⇄ 1−) would double the charging capacity of this anolyte and thereby lower the equivalent weight. Moreover, charging through the second, and more negative, redox couple would increase the OCV of the resulting RFB. However, to date, examples of multielectron Received: June 28, 2017 Accepted: September 12, 2017
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http://pubs.acs.org/journal/aelccp
Letter
ACS Energy Letters cycling of nonaqueous anolytes remain rare.15,16 Such experiments have been plagued by limitations in the stability and/or solubility of the tested anolytes, as well as by incomplete discharge through each redox event.16 A key objective of this study was to identify conditions for stable 2e− cycling of 1. Second, the supporting electrolyte contributes to the equivalent weight by providing the counterions to 1+ and 1−. As such, it is critical to replace the ammonium-based electrolyte (TBAPF6, 387 g/mol) with a lower-molecular weight and less expensive support. Previous literature reports have focused on the use of lightweight lithium-based supporting electrolytes. However, stable anolyte cycling in LiPF6 (152 g/mol) or LiBF4 (94 g/mol) is rare, primarily because of the low stability of reduced anolytes in the presence of these salts.17,18 These stability issues tend to be particularly pronounced for lowpotential anolytes like 1+.19−24 Furthermore, lithium cations are strongly Lewis acidic, and their coordination to reduced anolytes can detrimentally increase the redox potentials of these species. Such interactions have been exploited to increase the redox potentials of cathodic materials.25 However, an analogous increase in anolyte potential reduces the OCV, thereby negatively impacting energy density. As such, a second objective of this study was to identify a lightweight supporting electrolyte that enables stable 2e− cycling of 1+ without a significant attenuation of redox potential. In this Letter, we demonstrate that two-electron cycling of 1+ can be achieved using KPF6 as the supporting electrolyte. This support offers the advantage of dramatically lower molecular weight and cost relative to TBAPF6, while enabling stable cycling at low potential (−1.7 V vs Ag/Ag+). Ultimately, we demonstrate two-electron charge−discharge cycling of 100 mM solutions of 1+ in a flow cell. Using KPF6 as the support, this system undergoes 150 cycles with no detectable anolyte degradation with an equivalent weight of 166.2 g mol−1 e−1. One-Electron H-Cell Anolyte Cycling. We first evaluated the impact of supporting electrolyte (MPF6) on the first redox process of 1+ (1+ ⇄ 1). Galvanostatic charge−discharge experiments in a static H-cell were conducted using 5 mM solutions of 1+ in acetonitrile with 0.5 M of MPF6. A 2C cycling rate was used with a voltaic cutoff of −1.3 V vs Ag/Ag+. All experiments showed Coulombic efficiencies (CEs) between 96 and 98%. With TBAPF6 as the support, 1+ exhibits stable 1e− charge−discharge cycling, with less than 3% capacity loss over 50 cycles (Figures 1a and S1, black).14 In marked contrast, cycling in solutions of LiPF6 results in nearly 30% capacity loss over 50 cycles (Figures 1a and S1, green). We noted that the acetonitrile solutions of LiPF6 were mildly acidic,26 with pH values as low as 4.3. Literature reports suggest that this acid is formed via the disproportionation of LiPF6 to generate LiF and PF5 (eq 2).27,28 The PF5 can then react with traces of water to generate HF. This HF (or the electrophilic PF5 intermediate) is likely responsible for the capacity fade in this system, as it can react with the nucleophilic radical 1 that is formed upon reduction. 29 Thus, we hypothesized that removing traces of adventitious H2O and/or scavenging PF5 could enhance the 1e− cycling performance with LiPF6. Indeed, the addition of molecular sieves (to remove water) or N,Ndimethylacetamide (to scavenge PF5)30,31 did lead to some improvement in capacity retention (∼20% capacity loss over 50 cycles, Figure S4). Additionally, moving from LiPF6 to LiTFSI as the supporting electrolyte also improved the capacity retention (Figure S5). Overall these data suggest that moving
Figure 1. (a) Normalized discharge capacity versus cycle number for 1e− H-cell cycling (5.0 mM 1+, 0.50 M supporting electrolyte, 5 mA current). (b) Cyclic voltammetry (CV) of 1+ in various supporting electrolytes and the effect of electrolyte on the redox potentials (5.0 mM 1+, 0.50 M electrolyte, 100 mV/s scan rate).
away from Li salts will be advantageous for significantly enhancing cycling stability.
We hypothesized that the strong Lewis acidity of Li+ and the formation of the stable salt LiF both drive degradation of LiPF6 (eq 2). In contrast, TBA+ is a much weaker Lewis acid and has less driving force to abstract fluoride and form TBAF.32−35 We reasoned that Na+ and K+ should be more similar to TBA+, based on the fact that they are significantly less Lewis acidic than Li+ (ΔHhydration in kcal/mol: Li+ = −125; Na+ = −97; K+ = −77)36 and also have significantly larger cation radii (Li+ = 1.09 Å; Na+ = 1.50 Å; K+ = 1.97 Å).37 The use of NaPF6 resulted in only modest improvement of the 1e− cycling performance of 1+ relative to that in LiPF6 with 25% versus 30% capacity loss over 50 cycles (Figures 1a and S1, blue and green). However, the use of KPF6 led to a significant enhancement in cycling stability with only 5% capacity loss over 50 cycles (Figures 1a and S1, red).38 As illustrated in Figure 1b, CV of 1+ reveals that varying the cation of the supporting electrolyte from TBA+ to K+ to Na+ to Li+ has minimal impact on the reversibility or the standard potential of the first reduction event. In all cases, the potential (E11/2) is between −1.14 and −1.18 V. Despite these similarities by CV, the choice of cation has a profound impact on stability toward 1e− cycling. Most notably, KPF6 enables 1e− cycling of 1+/1 with capacity retention comparable to that with TBAPF6. In addition to good cycling performance, KPF6 offers the advantage of a much lower gravimetric mass than TBAPF6 (184 vs 387 g/mol). Furthermore, it imparts greater solution conductivity than analogous lithium salts (84 vs 70 S·cm2/ mol in acetonitrile, respectively).39 In addition, potassium is sourced from commodity chemicals that are over an order of 2431
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of the pyridinium cation 1+ to the radical 1 occurs at the first redox potential (−1.1 V). However, as the concentration of 1+ decreases and the concentration of 1 increases, it becomes kinetically more favorable to reduce 1 relative to 1+. The resulting enolate 1− then undergoes comproportionation with remaining 1+ in solution to form 1. This mass-transfer limited event results in charging at the potential of the second redox event (−1.7 V) before one electron per molecule has been charged (blue dashed line at 0.5 normalized capacity, Figure 2b). We note that this issue needs to be considered for all multielectron redox materials.16 While such a comproportionation does not diminish RFB capacity, it does lead to a decrease in energy efficiency. It can be addressed by either increasing the anolyte concentration or reducing the charging rate. Both result in suppression of this mass-transfer limitation and allow for almost all of the pyridinium 1+ to undergo reduction before the radical 1 is reduced. As an example, solutions containing 2.5, 5, and 10 mM anolyte 1+ charged to 0.6, 0.8, and 0.9 electrons per molecule, respectively, before 1 is reduced. Consequently, cycling of the 10 mM solution has an average energy efficiency of 96% compared to the 83% efficiency of the 2.5 mM solution. Flow Cell Cycling. We next transitioned from low concentration cycling in an H-cell to higher concentration cycling in flow. In addition to its direct relevance to RFB applications, flow cell cycling eliminates the large charging overpotential common to H-cell setups. Such overpotentials can lead to degradation pathways that are not representative of those occurring in RFBs.42 The 2e− flow cycling experiments were performed in a symmetric RFB using a Daramic microporous separator, carbon felt electrodes, and graphite charge-collecting plates with an interdigitated flow field (Figure S7).43,44 The reservoirs were symmetrically loaded with 100 mM of chemically synthesized14 radical 1 and 0.5 M KPF6 in acetonitrile. This concentration was selected for the experiment after the radical 1 was identified as the solubility-limiting redox state of the anolyte. Although an initial maximum solubility of 260 mM was measured for 1 in 0.5 M KPF6, this was found to be a kinetic solubility. Prolonged storage (14 days) of the 100 mM electrolyte solution led to a precipitation of 1, indicating that even the concentration used for this experiment was above the thermodynamic solubility limit. As summarized in Figure 3a, this setup enables cycling of 1 through all of its redox states, with the molecule serving as its own counter redox couple. Starting from the radical 1, reduction was performed at the working electrode with a current density of 20 mA/cm2 to form the enolate 1−. Simultaneously, the radical in the counter chamber was oxidized to the cation 1+. Under these conditions, the material in both reservoirs is continuously cycled through two electrons as opposite redox processes. This yields an energy-neutral flow cell with an average zero volt potential between the two electrodes, enabling the evaluation of electrochemical cycling and solution stability of 1+, 1, and 1− under continuous operation. Any decomposition or loss of one of the three components will result in a measured loss of capacity. Achieving stable cycling in such an experiment is a crucial requirement for obtaining a stable battery when the anolyte is ultimately paired with a catholyte. Examples of the charge−discharge plateaus for cycles 5−7 are illustrated in Figure 3b. The two plateaus at +0.6 and −0.6 V represent the two redox processes of the anolyte, referenced
magnitude less expensive than lithium commodities on a molar basis.40,41 Two-Electron H-Cell Anolyte Cycling. We next studied the impact of the supporting electrolyte on the 2e− cycling of 1+. CV studies show that the nature of the cation has a dramatic influence on the standard potential of the second redox couple (E21/2, 1/1−). As shown in Figure 1b, this second redox event occurs at the lowest potential with TBAPF6 as the support (E21/2 = −1.80 V versus Fc/Fc+). The potential then increases to E21/2 = −1.66, − 1.60, and −1.44 V in K+-, Na+-, and Li+-based electrolytes, respectively. This trend can be rationalized based on the relative Lewis acidities of the cations, which directly impact the strength of the interaction with the Lewis basic oxyanion 1−. Stronger interactions between 1− and the electrolyte cation stabilize the anionic anolyte and thus increase the reduction potential. Importantly, low potential redox events are critical for achieving high energy density anolytes. From this perspective, the TBA+ and K+ salts are preferred over the Li+ and Na+ analogues. Two-electron charge−discharge cycling between 1+/1/1− was next examined with KPF6 and TBAPF6 (poor 1e− cycling in LiPF6 and NaPF6 precluded stable 2e− cycling with these electrolytes). The experiments were conducted in a static H-cell using a procedure analogous to that in Figure 1a, except with voltaic cut offs of −1.8 V (for KPF6) and −1.9 V (for TBAPF6). Under these conditions, a 90% state-of-charge of 1− was reached, and similar capacity losses were observed for TBAPF6 and KPF6 [19% and 25%, respectively after 50 cycles, (Figures 2a and S2)]. CEs of 99% were observed for both experiments. The Nernst curves for these 2e− cycling experiments reveal a concentration-dependent asymmetry in the charge−discharge profiles of the two redox events (Figure 2b). Initially, reduction
Figure 2. (a) Discharge capacity versus cycle number for twoelectron charge−discharge bulk electrolysis in an H-cell of 2.5 mM 1+ in 0.5 M TBAPF6 (black) and KPF6 (red) electrolytes. (b) Normalized charge−discharge curves of the fifth cycle using 2.5, 5, and 10 mM 1+ at 5 mA charge−discharge rates. The dashed line indicates where 1e−/molecule is reached upon charging. 2432
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loss over 100 cycles). Furthermore, CV and 1H NMR spectra of the electrolyte after this four-day charge−discharge experiment were nearly indistinguishable from those of the initial solution (Figure 5 inset and Figure S10), indicating that there was no
Figure 5. Equivalent weight per electron transferred for 1+ with various MPF6 electrolytes for one-electron (red bar) and twoelectron transfer events (green bars). The blue dashed line represents the technoeconomic target.
degradation of the anolyte. These results demonstrate the high stability of the anolyte in all three redox states in MeCN/KPF6 and represent the first example of two-electron cycling of a low potential anolyte at high concentration with an alkali-metal supporting electrolyte. Figure 5 summarizes the equivalent weight of 1+ in the various electrolytes. This analysis reveals that by combining 1+ with KPF6 and simultaneously moving from one to two electron-transfer events, the equivalent weight is reduced from 312.8 to 166.2 g mol−1 e−1 (very close to the 150 g mol−1 e−1 target), while providing excellent cycling stability. This is in marked contrast to two-electron cycling with TBAPF6, which leads to an equivalent weight of 217.1 g mol−1 e−1. Although LiPF6 and NaPF6 electrolytes would, in principle, provide lower equivalent weights than KPF6, they do not support stable twoelectron cycling. It is also important to note that, because of the lower potential of the second redox couple of 1+, the theoretical energy density more than doubles upon moving from one to two electron-transfer events per molecule. Overall, these studies show that KPF6 provides the best balance of equivalent weight, energy density, cycling stability, and cost for the 2e− cycling of 1 +. In summary, this Letter describes a systematic evaluation of the effect of alkali-metal salts on the electrochemical cycling of the RFB anolyte 1+. We demonstrate that anolyte solutions supported by lithium- and sodium-based electrolytes rapidly lose capacity during 1e− and 2e− charge−discharge cycling in an H-cell. In contrast, KPF6 supports anolyte cycling that is comparable to that with the more expensive, higher molecular weight electrolyte TBAPF6. We further demonstrate that this stable two-electron cycling in KPF6 can be translated to higher concentration (100 mM) cycling in a flow cell. Under these conditions, 1+ was cycled through both redox couples for 150 charge−discharge cycles with no detectable decay of the anolyte. This system represents an equivalent weight of 166.8 g mol−1 e−1, which is close to technoeconomic targets for nonaqueous flow battery technologies. Overall, the exceptional cycling stability and compatibility of all three anolyte redox states (1+, 1, 1−) with a low molecular weight supporting electrolyte is promising for applications in nonaqueous RFBs. We also anticipate that the replacement of ammonium-based
Figure 3. (a) Diagram of the symmetric flow cell. (b) Charge− discharge curve of cycles 5−7.
against one another, and are consistent with the 500 mV separation of the redox couples measured by CV. The additional ∼100 mV ohmic drop primarily originates from membrane resistance, as determined by electrochemical impedance spectroscopy and polarization curve measurements (∼5 Ω·cm2, Figure S8).45 Figure 4 shows the charged and discharged capacities plotted as a function of cycle number. Using 10 mL of a 100 mM
Figure 4. Charge−discharge capacity versus cycle number of 100 mM anolyte in flow over two redox couples. The inset shows the CV of an aliquot of the electrolyte before and after 150 cycles.
solution of 1 (2.68 Ah/L theoretical capacity), the charge capacity reaches 23 mAh, thereby utilizing 86% of the available redox material. This is comparable to cycling 200 mM of a oneelectron anolyte. The small capacity loss over the first 50 cycles is due to initial equilibration of the flow setup (Figure S9). After this initial equilibration, the system was cycled over 100 times at 40 mA/cm2 without significant loss in capacity (0.28% 2433
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(11) Darling, R. M.; Gallagher, K. G.; Kowalski, J. A.; Ha, S.; Brushett, F. R. Pathways to low-cost electrochemical energy storage: a comparison of aqueous and nonaqueous flow batteries. Energy Environ. Sci. 2014, 7, 3459−3477. (12) Sevov, C. S.; Brooner, R. E. M.; Chénard, E.; Assary, R. S.; Moore, J. S.; Rodríguez-López, 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. (13) 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. (14) Sevov, C. S.; Hickey, D. P.; Cook, M. E.; Robinson, S. G.; Barnett, S.; Minteer, S. D.; Sigman, M. S.; Sanford, M. S. Physical Organic Approach to Persistent, Cyclable, Low-Potential Electrolytes for Flow Battery Applications. J. Am. Chem. Soc. 2017, 139, 2924− 2927. (15) Sevov, C. S.; Fisher, S. L.; Thompson, L. T.; Sanford, M. S. Mechanism-Based Development of a Low-Potential, Soluble, and Cyclable Multielectron Anolyte for Nonaqueous Redox Flow Batteries. J. Am. Chem. Soc. 2016, 138, 15378−15384. (16) Laramie, S. M.; Milshtein, J. D.; Breault, T. M.; Brushett, F. R.; Thompson, L. T. Performance and cost characteristics of multielectron transfer, common ion exchange non-aqueous redox flow batteries. J. Power Sources 2016, 327, 681−692. (17) Wei, X.; Xu, W.; Huang, J.; Zhang, L.; Walter, E.; Lawrence, C.; Vijayakumar, M.; Henderson, W. A.; Liu, T.; Cosimbescu, L.; et al. Radical Compatibility with Nonaqueous Electrolytes and Its Impact on an All-Organic Redox Flow Battery. Angew. Chem., Int. Ed. 2015, 54, 8684−8687. (18) Carino, E. V.; Diesendruck, C. E.; Moore, J. S.; Curtiss, L. A.; Assary, R. S.; Brushett, F. R. BF3-promoted electrochemical properties of quinoxaline in propylene carbonate. RSC Adv. 2015, 5, 18822− 18831. (19) Potash, R. A.; McKone, J. R.; Conte, S.; Abruña, H. D. On the Benefits of a Symmetric Redox Flow Battery. J. Electrochem. Soc. 2016, 163, A338−A344. (20) Duan, W.; Vemuri, R. S.; Milshtein, J. D.; Laramie, S.; Dmello, R. D.; Huang, J.; Zhang, L.; Hu, D.; Vijayakumar, M.; Wang, W.; et al. A symmetric organic-based nonaqueous redox flow battery and its state of charge diagnostics by FTIR. J. Mater. Chem. A 2016, 4, 5448− 5456. (21) Burgess, M.; Hernandez-Burgos, K.; Cheng, K. J.; Moore, J. S.; Rodriguez-Lopez, J. Impact of electrolyte composition on the reactivity of a redox active polymer studied through surface interrogation and ion-sensitive scanning electrochemical microscopy. Analyst 2016, 141, 3842−3850. (22) Park, S.-K.; Shim, J.; Yang, J.; Shin, K.-H.; Jin, C.-S.; Lee, B. S.; Lee, Y.-S.; Jeon, J.-D. Electrochemical properties of a non-aqueous redox battery with all-organic redox couples. Electrochem. Commun. 2015, 59, 68−71. (23) 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. (24) Wei, X.; Duan, W.; Huang, J.; Zhang, L.; Li, B.; Reed, D.; Xu, W.; Sprenkle, V.; Wang, W. A High-Current, Stable Nonaqueous Organic Redox Flow Battery. ACS Energy Lett. 2016, 1, 705−711. (25) Hernández-Burgos, K.; Rodríguez-Calero, G. G.; Zhou, W.; Burkhardt, S. E.; Abruña, H. D. Increasing the Gravimetric Energy Density of Organic Based Secondary Battery Cathodes Using Small Radius Cations (Li+ and Mg2+). J. Am. Chem. Soc. 2013, 135, 14532− 14535. (26) The pH of a 5% solution (v/v) of the 0.5 M LiPF6/acetonitrile electrolyte and water was measured. (27) Kawamura, T.; Okada, S.; Yamaki, J.-i. Decomposition reaction of LiPF6-based electrolytes for lithium ion cells. J. Power Sources 2006, 156, 547−554.
electrolytes with potassium analogues may prove more broadly applicable to other low-potential RFB anolytes. Our ongoing studies are focused on improving the solubility of 1 as well as pairing 1+/ KPF6 with a high potential catholyte and a suitable membrane in order to achieve a high energy density flow battery.
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acsenergylett.7b00559. Electrochemical protocols, cell designs, experimental details, and additional figures (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Melanie S. Sanford: 0000-0001-9342-9436 Author Contributions §
K.H.H. and C.S.S. contributed equally.
Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This research was supported by the Joint Center for Energy Storage Research (JCESR), a U.S. Department of Energy, Energy Innovation Hub.
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REFERENCES
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