Elucidating Factors Controlling Long-Term Stability of Radical Anions

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C: Energy Conversion and Storage; Energy and Charge Transport

Elucidating Factors Controlling Long-Term Stability of Radical Anions for Negative Charge Storage in Nonaqueous Redox Flow Batteries Jingjing Zhang, Jinhua Huang, Lily A Robertson, Rajeev S. Assary, Ilya A. Shkrob, and Lu Zhang J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b01434 • Publication Date (Web): 04 Apr 2018 Downloaded from http://pubs.acs.org on April 4, 2018

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The Journal of Physical Chemistry

Elucidating Factors Controlling Long-Term Stability of Radical Anions for Negative Charge Storage in Nonaqueous Redox Flow Batteries Jingjing Zhang, 1,2 Jinhua Huang, 1,2 Lily A. Robertson, 1,3 Rajeev S. Assary, 1,4 Ilya A. Shkrob, 1,2* and Lu Zhang 1,2 1

Joint Center for Energy Storage Research, Argonne National Laboratory, Argonne,

Illinois 60439, USA 2

Chemical Sciences and Engineering Division, Argonne National Laboratory, Argonne,

Illinois 60439, USA 3

Department of Chemistry, University of Illinois at Urbana-Champaign, 405 N. Mathews

Avenue, Urbana, Illinois, 61801, USA 4

Materials Science Division, Argonne National Laboratory, Argonne, Illinois 60439,

USA Corresponding authors: * Ilya A. Shkrob ([email protected]), Phone: 630-252-9516;

ABSTRACT Radical anions of electrochemically-reduced compounds (anolytes) have been suggested for storage of negative charge in nonaqueous redox flow batteries. The lower the redox potential of the anolyte molecule, the higher is the stored energy density. However, the stability of the radical ions frequently suffers as their redox potentials become extreme, and there is a compromise between the energy density and the chemical stability in the active form. In this study, we scrutinize this trade-off using one such “extreme,” the heterocyclic anolyte 2,1,3-benzothiadiazole, BzNSN, by adjusting the redox potential of BzNSN via installed electron-donating and electron-withdrawing groups. We show that the stability of the radical anion strongly depends on the degree of ion pairing in solution, with the worst being for the contact lithium ion pairs. For BzNSN derivatives, there is a strong correlation between the lifetime of the radical anion and the redox potential. The 1 ACS Paragon Plus Environment

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root cause appears to be the proton transfer from the solvent (acetonitrile) to the radical anion, which is concerted with the dimerization of the solvent molecules. These interactions cause the radical anion to become unstable even though the redox couple falls well within the electrochemical stability window of the solvent. Steering this redox potential towards the middle of this window using electron withdrawing groups did not pay off as it opened additional decomposition pathways. Our study, therefore, suggests that there can be natural limitations to the energy density that is realistically achieved using neutral, closed-shell anolyte molecules for charge storage in the redox flow cells.

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INTRODUCTION Anolytes (A) and catholytes (C) are redox active materials (ROMs) that are used as charge carriers in redox flow batteries (RFBs).

1-6

The energy density of cathodic and

anodic fluids in RFBs is defined by the difference in the redox potentials for the catholyte (that stores positive charge) and anolyte (that stores negative charge) in their respective compartments and the concentration of the charge carriers. The greater this potential difference, the more energy is stored, so that the most efficient battery utilizes as much of the electrochemical stability window (ESW) of the cell fluids as possible. In aqueous solutions, this window is limited due to electrolysis of water on the electrodes, but in solutions of aprotic polar organic molecules containing 0.1-2 M electrolyte salts and/or ionic liquids, this window can be several volts, which makes it potentially possible to reach much higher energy density than in the aqueous electrolytes for the same concentration and state-of-charge (SOC) in the device. 7-11 In practice, however, it proved difficult to profit from this potential advantage, for several reasons.

12-17

First, as the

redox couples move towards the extremes of the ESW, parasitic reactions make the charged states unstable, compromising the operation of the cell for which long-term stability of electrochemically separated charge carriers is required. Second, the solubility of ROMs in both states of charge can be too low to take advantage of the increased cell voltage. Third, as presently there are no selective separators for molecular ROMs in nonaqueous RFBs (NRFBs), crossover of these molecules and the products of their decomposition eventually occurs causing irretrievable capacity loss. All three of these complications are being presently addressed in the laboratories across the Joint Center for Energy Storage Research. In this study we focus on the stability of ROMs as it is the latter that ultimately defines what energy densities can be practically achieved in the NRFBs. In our previous studies, we sought to establish

18

what limits the stability of

catholytes using dialkoxybenzene derivatives as model compounds.

18-20

Here, we turn

our attention to anolyte molecules, once again limiting ourselves to a single class of such compounds: the derivatives of 2,1,3-benzothiadiazole (BzNSN), which is compound 1 shown in Table 1. Fully reversible one-electron reduction of BzNSN can be observed in 3 ACS Paragon Plus Environment


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acetonitrile, ethers, carbonates, and other polar organic electrolytes, and the resulting radical anion (1-●) is sufficiently stable to cycle the NRFB device. 21 To our knowledge, currently it is the only molecular design that allows sustained flow cell operation for an anolyte with the redox potential lower than -1.5 V vs Ag/Ag+ in acetonitrile, and 2.4 V operation of NRFB using this ROM was demonstrated when 1 was coupled with a dialkoxybenzene-based catholyte ROMs.

21

The use of acetonitrile in such NRFBs is

dictated by the necessity to reduce ion pairing in the practically important high concentration regime. Relatively weak interactions between the ions are needed in order to maintain high ionic conductivity. Only the most polar organic solvents allow using 0.5-2 M ROMs with the equivalent concentration of supporting electrolyte. Using derivatization of BzNSN with electron-donating or electron-withdrawing groups, one can vary the redox potential E0 from -1.68 V for 4, having a methoxy substituent, to 0.925 V for 8, having two cyano substituents (see Table 1; hereafter all potentials are given vs Ag/Ag+ if not specified otherwise). 22 This tunability allowed us to study how the proximity to the ESW edge affects the stability of the radical anion both in the electrolyte bulk and in the actual electrochemical cell. Since these anolyte ROMs presently have the lowest redox potentials, we aimed to learn from this example the general issues that emerge as one approaches the regime that alone justifies the use of NRFBs over their aqueous counterparts. Simplistically, one can expect that lowering of the redox potential will increase chemical reactivity of the radical anions, opening new pathways for unwanted parasitic reactions; i.e., some trade-offs between the energy density and chemical stability can be expected. Our study seeks to chart this balance for an important class of anolyte ROMs. A closely related issue is the role of electrolyte in such trade-offs. Charged anolyte molecules are radical anions (A-●) that can pair with cations (X+). The latter are always present in cell fluids in order to maintain high current density. Strong effects of ion pairing can be expected based on the reactivity of radical anions in the resulting contact pairs. These interactions depend on the nature of the cation, the radical anion, and the polarity of the solvent. We demonstrate that for BzNSN anolytes such ion pairing


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interactions overtake structural variations in controlling the chemical stability of the radical anions in acetonitrile-based electrolytes. The study proceeds as follows. Chemical reduction of 1 by alkali metals in ethers is used to demonstrate the formation of contact ion pairs in such solvents. Next, we show that the contact Li+A-● pairs also form in CH3CN and characterize the strength of interactions in different BzNSN derivatives. Then we characterize long-term stability of radical ions using magnetic resonance. Finally, we present product analyses aiming to establish the nature of reactions responsible for the radical ion stability. In a sister publication,

23

we will consider the performance of the same anolytes during

electrochemical cycling, where chemical interactions become more complex; in this study we are more concerned with the fundamental limits imposed by the redox system itself, as opposed to the complex cell environment. To save room, supporting tables and figures have been placed in the Supporting Information. When referenced in the text, these materials have the designator "S," as in Figure S1. METHODS Unless indicated otherwise, all chemicals were obtained from Sigma-Aldrich and used in their purest forms as received. Compound 5 in Table 1 was obtained from Matrix Scientific, other BzNSN derivatives were synthesized as described in refs. 22, 23 and 3 was synthesized as described in Section S1 in the Supporting Information. Nuclear magnetic resonance (NMR) spectra of 3 is given in Figures S1 and S2 therein. The chemical shifts (δ) are given in the units of parts per million (ppm) and the spin-spin coupling constants are given in Hz. Tetramethylsilane for 1H and

13

C and CFCl3 for

13

F were used as the

standards. Cyclic voltammetry (CV) data used in Table 1 were collected for 5 mM BzNSN solutions at 0.2 V/s using a CHI760D electrochemical workstation (CH Instruments, TX) in a three-electrode configuration with a Teflon encased glassy carbon disk working electrode, a Pt wire counter electrode, and an Ag/Ag+ reference electrode containing 10 mM AgNO3 in CH3CN. Electrolysis in the galvanostatic regime was conducted in the argon atmosphere using a borosilicate glass H cell equipped with a ceramic porous 5 ACS Paragon Plus Environment


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separator; the cell fluids were agitated by a magnetic stirring rod during the electrolysis. Each cell compartment initially contained ≈5 mL of the same solution to minimize ROM crossover. Reticulated vitreous graphite electrodes were used in both of the cell chambers, and the solutions was charged to 100% SOC at a rate of 5C, which corresponds to charging to the full capacity in 1/5 h. Parstat MC potentiostat (Princeton Applied Research) was used to control the H cells. The working chamber of the cell was equipped with a Ag/Ag+ reference electrode to control voltage. Typically, to reduce BzNSN the voltage on the carbon cathode was controlled from -2.4 V to -0.5 V, and the cell charging time was 10 min. In some experiments, voltage control of catholyte oxidation was used instead. Electrochemically-generated radical ions in the harvested cell fluids were observed

using

continuous-wave

electron

paramagnetic

resonance

(cw

EPR)

spectroscopy in the X-band using 100 kHz field modulation. The details of these measurements are given in Section S2 in the Supporting Information. Decay kinetics were obtained using the same EPR spectrometer at 25 °C. After the complete decay of the paramagnetic species, the harvested cathodic and anodic cell fluids were analyzed chemically. Note than in some cases, this waiting period extended to 2+ months. 1H, 19F and

13

C NMR spectra were obtained using a Bruker Avance III HD spectrometer (300

MHz), and the same instrument was used to perform 1H-1H correlation spectroscopy (COSY) and diffusion-ordered NMR spectroscopy (DOSY) that was used to separate signals in complex reaction mixtures. Fourier transfer infrared (IR) spectra were obtained using a Nicolet iS5 spectrometer (ThermoFisher Scientific) inside a glove box. The powder was pressed against a thin silicon disk. We also analyzed the charged cell fluids using gas chromatography – mass spectrometry (GC-MS) and high performance liquid chromatography (HPLC) using the same instrumentation as in our recent studies, see refs.

18, 20

For GC-MS, additional

preparation was required to remove nonvolatile salts. To this end, the solvent was removed in vacuum, and water and CH2Cl2 were added. The extracted material from the dried organic layer was either analyzed directly or separated further by elution with CH2Cl2 or CH2Cl2: CH3CN (5:1 v/v) on short Hypersep silica gel column cartridges. The 6 ACS Paragon Plus Environment

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column residue was subsequently stripped by acetonitrile, and the eluted fractions were concentrated and analyzed separately using NMR, HPLC, and GC-MS. To calculate proton hyperfine coupling constants (hfcc’s), density functional theory (DFT) calculations using B3LYP functional 26

implemented in Gaussian 09

24, 25

with 6-31G(d,p) basis set as

was used. Geometry optimization and gas-phase

energetics were calculated using the same method, while the redox potential was estimated as described in the literature. 18, 27 The effect of the solvent was simulated using the polarized continuum model, as explained by Assary et al.

27

All reactions energetics

were estimated using the latter model. RESULTS AND DISCUSSION Chemical reduction of 1. Following Atherton,

28

alkali metals were used to reduce 1 in 1,2-

dimethoxyethane, DME (see section S2 for more experimental detail). Metallic sodium or potassium was placed in a vial containing 10 mM 1 in DME to generate 39

23

Na+1-● and

K+1-● pairs, respectively. Naturally abundant alkali nuclei are magnetic with the spin of

3/2, and one may observe isotropic hfcc on these nuclei provided that there is non-zero unpaired electron density in the corresponding s-orbitals. To obtain 7Li+1-● pairs, 0.1 M lithium bistriflimide (LiTFSI) was added to the DME solution containing 10 mM Na+1-● pairs: at this concentration, lithium ions completely substituted sodium ions in the contact pairs. The resulting EPR spectra are shown in Figure 1, and Table 2 summarizes hfcc’s obtained by simulation of these EPR spectra. In all three pairs, non-zero electron density in the spin bearing s-orbital of the Alk+ ion is observed. This unpaired electron density systematically decreases with the increasing ion radius, suggesting the formation of a close contact pair. Concomitantly, the hfcc’s in nitrogen-14 decrease in the direction from Li+ to K+, while the 2,5-proton hfcc’s (see Tables 1 and 2) increase in the same direction. By adding 0.1 M tetraethylammonium bistriflimide (NEt4TFSI) to 10 mM Na+1-● solution from which unreacted sodium metal was removed, a complex EPR spectrum (Figure S3) was obtained indicating the formation of weakly interacting NEt4+1-● pairs pairs (64.5 mol%) 7 ACS Paragon Plus Environment


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in dynamic equilibrium with the contact Na+1-● pairs (35.5 mol%). Table 2 gives the hfcc constants for both of these species. The hfcc’s in the weakly interacting radical anion follow the trend observed with the increase in the ion radii. Further, comparison of hfcc’s for free and paired radical anions evidences that the alkali cation pulls the electron density from π−orbitals in the arene ring towards the nitrogen atoms that strongly interact with the alkali cation. The influence of ion pairing can also be observed electrochemically using CV (Figure S4). The redox potential of 1 is solvent dependent, being higher in more polar solvents (Figure S4a). In CH3CN, the redox potential decreases when NEt4TFSI replaces LiTFSI as salt, i.e., stronger ion pairing causes this redox potential to increase (Figure S4b). This trend is also evident in Figure S4c, where CV scans for the contact ion pairs involving Li+, Na+, and K+ ions are compared: the redox potential increases as the cation radius decreases. Thus, both EPR spectrometry and cyclic voltammetry indicate strong ion pairing between the smaller alkali ions and 1-●. When DME solution of Na+1-● was diluted 1:1 v/v with a polar solvent, tetrahydrofuran (THF), 20% of these contact ion pairs dissociated, (Figure S3 and Table 2) such that both paired and free radical anions were observed. Nevertheless, most of the ion pairs remained in contact even in such polar solution, suggesting strong electrostatic interactions in the contact pairs. Below we will show that contact Li+1-● pairs can also be observed in CH3CN. Figure S5 shows the decay kinetics for Na+1-● pairs in DME observed using EPR spectroscopy. The same kinetics can be followed using the optical absorption of 1-● that is responsible for the dark green coloration of the solution. The radical anion persists over 100 h, but eventually the EPR signal decays and the color fades. After this decay, extractive workup and product chemical analyses (see details in Section S2 and Table S1 and Figures S6 to S12 therein) gave different results depending on whether 1-● remained in contact with metallic sodium during aging of the solution. Without sodium, ophenylenediamine (P1 in Scheme 1) was the main product, whereas in the presence of Na0, there was also trans-o,o’-diaminodiazobenzene (P2 in Scheme 1), indicating ringopening and the loss of sulfur. Reductive sulfur extrusion and proton transfer 8 ACS Paragon Plus Environment

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The loss of sulfur during chemical reduction is known as reductive sulfur extrusion; it occurs for many of S-containing heterocycles, including BzNSN.

29, 30

For

BzNSN, many of the common organosynthetic reducing agents (e.g., Zn0 in acetic acid, lithium aluminum hydride, sodium borohydride, etc.) are known to reduce BzNSN to P1. 30

The exact mechanism for these reactions is unknown, but the symmetric 1,3-dihydro

intermediate (P3 in Scheme 1) has been postulated.

30

P3 can be synthesized by reacting

P1 with piperidine-1-sulphenylchloride in ether; this compound is unstable in chloroform and acetonitrile, yielding polymeric products of unknown composition.

31

Importantly,

there is evidence that in solution P3 is in the dynamic equilibrium with thionitroso tautomer P4 (Scheme 1). Such R-N=S compounds are known to dimerize, losing one or two sulfur atoms to yield R-N=S=N-R (that hydrolyze to RNH2 and SO2 when contacted with water) and R-N=N-R compounds.

32

The formation of P1 and P2 in DME (see

above) can be qualitatively explained by (i) protonation of 1-● at the nitrogen atom to yield the N-centered 1H● radical (Scheme 1), (ii) disproportionation of these radicals to yield P3 in the equilibrium with P4, and (iii) subsequent sulfur extrusion from the products of P4 dimerization. 32

P1

P2

P3

Alk+

P4

1H

P5

Scheme 1. Some Reaction Products of Reductive Sulfur Extrusion. Our DFT estimates suggest plausibility of this scenario (see Table S2 in section S3). In the gas phase, the protonation of 1-● at a nitrogen atom is preferred over any other 9 ACS Paragon Plus Environment


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position. The recombination of the N-centered 1H● radicals is endergonic by 0.17 eV, whereas their disproportionation is exergonic by 0.65 eV. For other derivatives in Table 1, this radical disproportionation is also exergonic (> 0.47 eV) but the energy gain for 8 and 9 is relatively small (0.16 eV and 0.28 eV, respectively; see Table S2). The N-H bond energy in P3 and P4 is only 3.03 eV and 3.15 eV (for 1), respectively (vs > 4 eV for the typical C-H bond), and the 1H● radical cannot abstract hydrogen from other molecules in the solution. It also cannot recombine and/or add to a double/triple bond, so these radicals may only decay by disproportionation to recover 1 and also yield P3 and/or P4. According to our DFT calculations, in the gas phase P4 is more stable than P3, having ≈ 0.12 eV lower energy. This stability difference is also the case for other compounds in Table 1 with an exception of 9: the energy gap between the two tautomers varies from 0.11 eV for 6 to 0.26 eV for 4, with the thionitroso form energetically preferred (Table S2). We believe that this tautomer is also prevalent in solution as DFT calculations in various polarized continuum solvation models yield the same energy ordering although the energy difference is small. Furthermore, our DFT calculations suggest that dimerization of P4 with the sulfur loss would be strongly exergonic (by 0.63 eV for R-N=S=N-R and 1.83 eV for R-N=N-R), and these reactions remain exergonic for all derivatives in Table 1. The branching between these reactions may depend on the form in which the sulfur is extruded, which in turn depends on the presence of Na0 in contact with the solution (the sulfur can form sulfides, changing the reaction energetics); without Na0, the nominal 4e-, 4H+ reduction to P1 seem to be the main reaction path, as in other examples of reductive sulfur extrusion found in the literature. 29, 30 While such a product is expected, it is puzzling where the protons for this reaction come from, since protogenic impurity should be eliminated by reaction with an alkali metal. It is logical to assume, therefore, that the solvent itself becomes deprotonated. As acetonitrile was used in our electrochemical experiments, we considered whether deprotonation of this solvent by 1-● is thermodynamically possible. Aliphatic nitriles are known as aprotic solvents; however, strong bases (e.g., NaNH2) can reduce these nitriles to amidines; for CH3CN the main product of this proton transfer is 310 ACS Paragon Plus Environment

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(sodioamino)crotonitrile (P5 in Scheme 1), which is surprisingly stable in organic solvents. Hydrolysis of P5 yields 3-aminocrotonitrile that is also known as diacetonitrile. 33, 34

In addition to the dimer compounds, trimer compounds (salts of 4-amino-2,6-

dimethylpyrimidine and 2,4,6-trimethy-1,3,5-traizine) are known to form under certain conditions,

35

including radiolysis of CH3CN,

36

and higher polymers can also form. The

dimers and trimers of CH3CN are the main products of electrochemical reduction of CH3CN on the Pt electrode that occurs around -3.3 V (i.e., the anolytic edge of the ESW), 35, 37, 38

which is much lower than the redox potentials for BzNSN derivatives in Table 1.

This electrochemical reaction is believed to involve the hydride anion generated by the reduction of chemisorbed hydrogen gas on the platinum surface. 35 While the formation of a bare NCCH2- anion would be prohibitive energetically, the dimerization is a concerted reaction in which this anion (unobserved as a separate species in CH3CN) is coupled to another CH3CN molecule. This concerted reaction considerably changes protonation energetics. To estimate the thermochemistry of the protonation of 1-● by CH3CN, we considered (CH3CN)3Li+ 1-● pairs shown in Figure 2 in which the tetrahedrally coordinated lithium cation was solvated by the nitrile groups of three CH3CN molecules and the anion; the same solvation mode was assumed for P5. The corresponding reaction of concerted deprotonation to yield 1H● was exothermic by 77 meV in the gas phase and endothermic by 100 meV in our utilized polarized continuum model.

39, 40

Thus, through a concerted reaction that involves contact Li+1-●

pairs, it is possible for the radical anion of 1 to deprotonate CH3CN in solution although such a reaction is weakly endergonic, i.e., relatively slow. Since the gas phase proton affinity of radical anions correlates linearly with the electron affinity of their parent compounds (see Figure S13) and the latter correlates with their redox potential, one can expect that the exothermicity and the rate of the protonation also correlates with the redox potential. This redox potential can be determined computationally using the method discussed in refs.

39, 40

See Figure S14 for correlation

of the computed and experimentally determined redox potentials. In Figure S15 we show the Hammett plot for a wider selection of 5-substituted BzNSN derivatives. This figure illustrates a good correlation between the computed redox potentials and the Hammett 11 ACS Paragon Plus Environment


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constant σm, which empirically takes into account the electron-withdrawing/donating properties of the substitutuent group. These plots allow semiquantitative prediction of the redox potentials and qualitative prediction of the radical anion stability provided that the reaction is the protonation. An important insight is that the energetics of the protonation should be very sensitive to the nature of the cation as strong Coulomb interactions both in the contact ion pair and the salt like reaction product are essential for favorable reaction energetics. With this insight, we proceed to examine the behavior of electrochemicallygenerated BzNSN radical anions in CH3CN solutions. Lithium ion pairing of BzNSN-● in CH3CN To study radical anions of BzNSN derivatives in CH3CN, electrochemical reduction of the parent compound was used. To this end, 5-50 mM solutions were placed in a bulk electrolysis setup and charged to 100% SOC. We found that 9-● decayed even during CV runs in CH3CN, while 6-● had such a short lifetime that we were unable to detect any EPR signals after electrolysis. For these reasons, we excluded 6 and 9 from further consideration. All other BzNSN derivatives yielded EPR spectra when the parent compounds were reduced in CH3CN. Figure 3 shows examples of the EPR spectra for 1-● and 2-● obtained in 0.5 M LiTFSI and 0.5 M tetrabutylammonium hexafluorophosphate (NBu4PF6) solutions. Different EPR spectral patterns were observed for the same radical anion in different electrolytes. When the supporting salt anions are exchanged (i.e., when LiPF6 and NBu4TFSI were used instead of LiTFSI and NBu4PF6, respectively), the resultant EPR signal patterns were the same as in Figure 3, revealing that the interaction of the radical anions with the cations determines the EPR pattern. Table 3 summarizes EPR parameters observed in the LiTFSI solutions, and Table 4 gives these parameters for the NBu4PF6 solutions. The same hfcc parameters were obtained in solutions containing other lithium and tetrabutylammonium salts, further supporting the fact that the supporting anion is mainly a spectator. Table 3 shows that non-zero unpaired electron density was observed on lithium-7 nuclei (7Li+), suggesting the formation of contact Li+ BzNSN-● pairs in acetonitrile despite its high polarity. By EPR of dilute electrolyzed solutions, we estimated that the association constant is > 103 M-1. Using a. c. conductivity 12 ACS Paragon Plus Environment

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measurements, we estimated that this constant is ~ (1-2)x104 M-1, so the ion pair remains strongly associated under the practically important conditions. There are other recent examples of strong association of the anolyte radical anions with Li+ and Mg2+ ions in such systems, see refs. 41, 42. Using the “minimalistic” DFT model of these contact ion pairs shown in Figure 2, in Tables S4 and S5 we optimized the geometry of this species and calculated hfcc parameters for free and paired radical anions that are given in Tables 3 and 4. These calculations yield close estimates for hfcc in 7Li+ ions, suggesting that the cation is located 3.51-3.52 Å away from the sulfur in BzNSN-●. When electron-withdrawing groups are introduced, this distance increases over 3.55 Å, and the spin density decreases considerably (Table S5); this trend is also seen in the experimental estimates for hfcc’s on lithium-7 (Table 3). Concomitantly, the spin density in both of the nitrogens decreases. This can be seen both in the isolated radical anions and in the contact ion pairs, by comparing Tables 3 and S5 and Tables 4 and S4, respectively. In Figure S16, we compare the computed and experimental hfcc’s. Overall, the agreement is excellent. Significant differences between hfcc’s in LiTFSI and NBu4PF6 solutions suggest strong Li+ ion pairing in the former solutions. In the NBu4+ solutions, the hfcc’s are much closer to the ones calculated using DFT for isolated radical anions. To summarize, the combined computational data coupled with EPR spectroscopy indicates that (i) Li+ ions strongly interact with BzNSN-● anions in CH3CN, changing electron density on the radical anions, (ii) larger organic cations interact weakly with these radical anions, (iii) the strength of ion pairing for Li+ decreases for more positive redox potentials that reflects the effect of the electron-withdrawing groups (see Figure S15) on the electron density in the nitrogen atoms, which is observed using EPR spectroscopy. Chemical stability of BzNSN-● anions Given that the cation exerts such a strong effect on the radical anion, even in polar solvents, we hypothesized that chemical reactivity follows the energetics and becomes



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different for strongly and weakly interacting cations. Below, we examine kinetic experiments suggesting that this is indeed the case.

10

DABCO

Scheme 2. Structural Formulas for Catholyte 10 and Sacrificial Electron Donor DABCO. For these studies, 50 mM solutions of BzNSN derivatives in CH3CN containing 0.5 M salt were used, and the decay of radical anions was followed using EPR spectroscopy. We found that at 1-5 mM anolyte (at which concentration the EPR line broadening is minimal, so it was used for the spectroscopic characterization), the decay kinetics of A-● were short and poorly reproducible from sample to sample; however, at 50 mM, after the initial fast decay these kinetics became long and reproducible (see, e.g., Figure 4). This behavior suggests the presence of a reactive impurity in the electrolyte; when the concentration of A-● exceeds the concentration of this impurity, after the initial rapid decay due to reaction with this impurity the remaining radical anions decay much slower. Importantly, as the anolyte molecules reduce during electrolysis in the working electrode chamber, another compound needs to oxidize in the counter electrode chamber. To reduce crossover of reaction products, the same concentration of ROMs was used in both of the cell compartments. If no catholyte is present, BzNSN becomes oxidized itself (see below). In this process, protons are released, and they can become transferred between the two compartments following the field lines. 23 The crossover of the oxidation products, including carbocations and protons, into the working electrode chamber can shorten A-● lifetime through inadvertent reactions. We believe that such crossover was 14 ACS Paragon Plus Environment

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minimal during 5C charging as no oxidation products were detected in the cathodic fluid. To eliminate this possibility entirely, we added catholytic ROMs such as ferrocene, which can be reversibly oxidized without deprotonation, or sacrificial electron donors that

serve both

as

oxidation

targets

and

proton

acceptors,

such

as

1,4-

diazabicyclo[2.2.2]octane, DABCO (Scheme 2). These additives we introduced into one or both compartments of the cell. In no instance did we observe the effect of such additives on the radical anion lifetime. Also, in the current NRFBs both the anolyte and catholyte molecules are added to the cell fluids in both compartments in order to minimize crossover. Therefore, radical anion stability in the presence of catholyte molecules is more important practically than the stability of an isolated radical anion. For this reason we used our standard dialkoxyarene catholyte 10 (1:1 equiv., see Scheme 2) in both cell compartments and compared the radical anion lifetimes with and without the catholyte molecules and found the catholyte molecules had no visible effect on A-● lifetime (e.g., Figure 4a). Furthermore, the A-● lifetime did not change significantly in the electrolyte solutions (at the same salt concentration) that contained salts with the same cation but different anions as illustrated in Figure 4b for 1-● in LiTFSI and LiPF6. In both cases, the decay kinetics were nearly exponential, and time constants t1/2 for 50% decay were ≈ 45 h vs ≈ 52 h, respectively. The most striking effect is observed when Li+ is replaced with a larger NBu4+ cation: the corresponding lifetimes become ≈ 709 h for NBu4TFSI and ≈ 853 h for NBu4PF6 (Figure S17). These observations suggest that strong ion pairing significantly decreases BzNSN -● stability in CH3CN. In Figure 5, we plot the experimental lifetimes for selected BzNSN-● anions paired with Li+ or NBu4+ cations as a function of the redox potential of BzNSN molecules. The increase in the radical anion stability by changing a small Li+ ion to a large organic cation is seen not only for 1, but also for 2, 3, and 4 in Table 1 (see Figure 6). Equally remarkable is the linear correlation of log t1/2 and the redox potential of the BzNSN molecules: as the electron affinity increases, and the anolyte molecule becomes reduced at a higher potential, the stability of the corresponding radical anion increases significantly. High as this increase can be, it is, actually, smaller than the effect of ion pairing (Figure 5). Interestingly, the lifetime of 10+● in the anodic fluid changes only slightly as the salt is varied (see Figures 5 and S17); i.e., the effect of the salt cation on 15 ACS Paragon Plus Environment


Page 16 of 42

the lifetime of an anolyte radical anion does not have a correlation to the effect of the salt anion on the lifetime of a catholyte radical anion. In Li+ solutions, the decay of 10+● is longer than the decay of all radical anions shown in Figure 5, while in NBu4+ solutions, the decay of 1-● and 2-● becomes longer than the decay of 10+●. As most NRFBs use Li+based electrolytes, the detrimental effect of Li+ ions on the anolyte stability in the contact pairs has the obvious practical implications. From the latter standpoint, the redox potential of the anolyte needs to be as low as possible to maximize electric energy density. In this regard, it is troubling that radical anions derivatized with ideally beneficial electron-donating groups (i.e., to lower the redox potentials below -1.6 V) are so markedly destabilized. Simple deduction predicts that encumbering the anolyte structure with the electron-withdrawing groups should increase the redox potential (Table 1) and result in more stable BzNSN-● radical anions, especially when using large organic cations such as NBu4+ to weaken the ion pairing. However, these expectations were not met. As previously described, 6-● with a trifluoromethane substituent was unstable in all electrolytes with a lifetime of just a few minutes. Fluorinated 5-● was stable in Li+ solutions but very unstable in NBu4+ solutions (Figure 7a), which is the opposite behavior of electron-donating homologs 1-4 shown in Figures 6a and 6b. Radical anion 7-● with an acetyl substituent was unstable in all electrolytes (Figure 7b). The only relatively stable BzNSN-● radical anion with a higher redox potential was 8-● containing two bridging cyano substituents (Figure 7b), for which E0(8/8-●) ≈ -0.93 V. However, it can be difficult to take advantage of this greater stability in practice as there is second potential E0(8-●/82-) ≈ -1.52 V for this anolyte molecule that is corresponding to the reversible 2e- reduction.

23

As the electrochemical cell becomes

cycled and the overpotential increases near the end of the charging cycle, it becomes difficult to avoid this second electron reduction. 23 Given these results, it becomes clear that the inclusion of the electronwithdrawing groups on the BzNSN molecule (as opposed to inclusion of the electrondonating groups) opens new reaction pathways for decomposition of the radical anion thereby shortening its lifetime. A good example of this trend is 7-●. It is known that radical anions of esters can undergo C-O bond scission 43 16 ACS Paragon Plus Environment

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R’C(O)OR + e- → R’CO2- + R●

(1)

For 7-●, most of the negative charge is in the π-system of the arene moiety; however, some of this charge is shared by the acetyl group; this is sufficient for the slow elimination of the methyl radical to occur. This side reaction can be demonstrated by chemical analyses of the solution obtained after decomposition of 7-● (Figures S18 to S20). The resulting product has the optical absorption (Figure S18) and 1H NMR (Figure S19) spectra that are very similar to 7, and in the gas chromatogram there is m/z 181 mass peak with the fragmentation pattern that is consistent with protonated 5-carboxy-BzNSN, which is the expected product of reaction 1. The lithium salt of this anion is poorly soluble in CH3CN. This solid residue is soluble in water, and the electrospray ionization (ESI) MS spectrum of the aqueous solution shows an anion with m/z -179.2 that also corresponds to the product of reaction 1. This telling example shows that introducing electron-withdrawing groups into the anolyte molecules aiming to increase their lifetimes can backfire due to opening of new reaction pathways and further compromise the performance due to the formation of poorly soluble products. Chemical analyses of anodic fluids Bulk electrolysis experiments are typically conducted by controlling the voltage and the current in one of the cell compartments using a reference electrode. In our case, the reference Ag/Ag+ electrode was placed in the cathode compartment to control anolyte reduction. In other words, the oxidation process on the anode was not controlled. As mentioned above, in many if not most of the ROM development studies, the catholyte and anolyte molecules are tested separately. We have found, however, that the cycling stability of the ROMs in one cell compartment was considerably influenced by the reactions in another cell compartment. Due to imperfect separators, the reaction products eventually cross over between the compartments and the chemical evolution becomes complex.

23

Here we briefly summarize the processes occurring in the counter electrode

chamber during our electrolysis experiments. A photograph in Figure S21 shows the vials containing the anodic and cathodic fluids with the intensely colored radical ions of the ROMs molecules.



Page 18 of 42

In the absence of a catholyte or other readily oxidizable additives, BzNSN become oxidized, yielding numerous products. In the bistriflimide anion containing solutions, the main product of this oxidation was the TFSI- adduct of BzNSN in the arene ring (product C in Figure S22). For most of the BzNSN derivatives, the 4-position adduct is prevalent, but for 6, the 7-position adduct is prevalent. Using isomers can be clearly distinguished for 5 due to non-zero

19

19

F NMR, the two

F-14N coupling observed

only in the 4-adduct. These adducts can also be observed using GC-MS, yielding the characteristic m/z peaks. It appears, therefore, that the TFSI- anion (>N-) is oxidized on the anode to the imidyl radical (>N●) that adds to the arene ring of the BzNSN molecule. When less oxidizable anions were used (e.g., PF6-) no adducts involving the fragments of these anions were observed, and the prevalent products were ring-to-ring dimers and pigments obtained by fusion of two BzNSN molecules through their NSN rings (products B and E in Figure S22). The exact manner of this fusion remains unclear. In addition to these species, several other minor products were found using chromatographic fractionation and 1H NMR spectroscopy (Figure S22). For some BzNSN molecules (e.g., 4) oxidation in their side groups was observed, too. None of these many side products was observed when a sacrificial electron donor DABCO (Scheme 2) was introduced into the anodic fluid. These products of BzNSN oxidation were also lacking when 10 was introduced; all products observed in the anodic fluid were derived from 10+●. There was no indication of oxidation of the electrolyte. Among the decomposition products indicated by GC-MS and 1H NMR were the product of tert-butyl radical loss from the arene group of 10 (both the monomer and the dimer were observed, see Figure S23), which are the main reaction products in the anodic fluids (see Figure S23). Also observed are the quinone, the products of the methyl group loss and addition, two transalkylation products (Figure S23) and other minor products (see the typical HPLC chromatogram in Figure S24). The tert-butyl carbocation eliminated from 10+● is observed as N-tert-butylacetamide using GC-MS. The formation of this product suggests direct involvement of the solvent as acceptor to yield t-BuN=C+Me cation whose hydrolysis yields the detected product. Identical reaction products for 10+● were observed in the presence of different BzNSN derivatives; i.e., no evidence for reactions of 10+● with the anolyte molecules was found. 18 ACS Paragon Plus Environment

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Chemical analyses of cathodic fluids To examine the chemistry of cathodic fluids, we charged the electrochemical cells to 100% SOC and waited for the complete decay of the radical anions before analyzing the composition of these solutions. We remind that NMR spectroscopy of “live” solutions containing radical ions is complicated by paramagnetic broadening of the resonance lines and rapid degenerate electron exchange between the radical anions with their parent molecules, which make the resonance lines “invisible” to NMR. When dilute solutions (5 mM) of BzNSN molecules were used in these experiments, the most recognizable reaction product was the corresponding diamine (P1 in Scheme 1); e.g., see Figure S24. In addition, in some samples there was a product with an ABCD proton system (see Figure S25) with the intact arene ring and asymmetric substitution, which can be product P4 in Scheme 1. This intermediate product decomposed during chromatographic analyses yielding the diamine, and it was observed by NMR only. The same products were observed at higher concentration of the anolytes, but the reaction yield was relatively small, suggesting (in accordance with our kinetic measurements) that protonation of BzNSN-● at low concentration involved protic impurities in the solvent. Instead, in the HPLC chromatograms, a wide, tailing peak corresponding to a polymer species absorbing at 250 nm was observed (Figure S26). This species was not observed in GC-MS and NMR. In LiTFSI solutions, a white precipitate formed that was not observed either in LiPF6 or NBu4TFSI solutions. This precipitate did not dissolve in organic solvents so it can be readily separated and washed from the traces of other products in the reaction mixture. We have obtained transmission infrared (IR) spectra of this material for different BzNSN molecules (Figure S27a); all of these IR spectra were similar to each other, suggesting the same composition; this excludes BzNSN molecules as the progenitor of the solid material. Comparison with the IR spectra of the ROMs and the salts also suggests that the solid residue is derived from the solvent alone. This solid material reacts and dissolves in water; neither 1H nor

19

F resonances

were observed in D2O, and the only signals identifiable by NMR were from the 7Li+ ions, indicating that this solid material is a lithium salt. GC-MS analysis of the aqueous 19 ACS Paragon Plus Environment


solutions

yields

3-aminocrotononitrile,

suggesting

that

Page 20 of 42

this

salt

can

be

3-

(lithioamino)crotonitrile that hydrolyzes in water; however, it can be another form of polymeric CH3CN that yields 3-aminocrotononitrile when hydrolyzed. This is consistent with our hypothesis that A-● decays in the electrolyte by deprotonating CH3CN solvent; it is difficult, however, to ascribe specific chemical structures to the products of this reaction, as they rapidly undergo further reactions, some of which involve polymerization. The broad HPLC peaks were observed for all of the studied BzNSN derivatives; whether the reaction resulted in the formation of insoluble residues depended on the electrolyte composition and BzNSN derivatization. In no case did we observe 1H and 13C NMR resonances from this polymer, indicating its rigidity and lack of chain dynamics that causes poor averaging of magnetic anisotropies. The lowest yield of this polymer was observed in the 1/ferrocene/LiTFSI system, in which only 610% of 1-● converted to P1. In Figure S28, the kinetics of the proton NMR signal from P1 is shown for electrochemical reduction of 50 mM 1 in CD3CN. The signal from 1-● cannot be observed due to the degenerate electron exchange, but the protons in P1 are not involved in this exchange, so P2 formation can be followed using 1H NMR signals from the arene protons. Both the radical anion decay and the P1 formation are exponential, and the corresponding t1/2’s are 52.2 h and 60.3 h, respectively. This close correspondence suggests that the rate limiting step in the stepwise formation of P1 is the decay of the radical anion. The yield of this product, however, was ~10%, and the main reaction product according to HPLC of the reaction mixture was the polymeric anionic compound. To put these analyses in perspective, while the reactivity of the anolyte radical ion strongly correlates with the redox potential (Figure 5), our chemical analyses as well as the prohibitive reaction energetics of the electron transfer do not suggest that the solvent reduction occurs in these electrolytes directly. As we argued above, the proton transfer energetics would also follow the order of the redox potentials (Figure S13). In CH3CN, concerted proton transfer and dimerization are only slightly endergonic. Protonated products derived from the anolyte molecules and diacetonitrile are both found among the reaction products. The formation of such reaction products is supportive of the stepwise 20 ACS Paragon Plus Environment

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reductive sulfur extrusion reactions discussed above. As in other instances of such sulfur extrusion, our understanding of such reactions remains incomplete. In particular, we cannot establish the structure of the polyanion product of the reaction, which is sadly also the case for many other S extrusion reactions. In our case, however, we are confident that the polymeric product is composed mainly of decomposed electrolyte. CONCLUSION In this study, we scrutinize the nature and reactions for electrochemically reduced high-energy anolyte molecules (BzNSN derivatives), which are used as negative charge carriers for long-term charge storage in NFRBs.

21

Currently, 1 has the lowest redox

potential among the anolyte molecules that sustain relatively stable operation of the allorganic NRFBs. As this redox potential can be changed through substitution in the anolyte molecule, our initial intent was to learn (i) whether this redox potential can be lowered without compromising the chemical stability of radical anion and (ii) how far can this stability be extended by tuning of the structure (regardless of the redox potential). Our study provides nuanced answers to both of these questions. The answer to the first question is that lowering of the redox potential through derivatizaton with the electron-donating groups always destabilizes the radical anion. We trace this destabilization to the increased proton affinity of these radical anions (which correlates with their redox potential), but the root problem is the favorable energetics of the proton transfer itself. While the direct proton transfer from the CH3CN molecule is prohibitive energetically, it becomes more favorable when this proton transfer becomes concerted with the dimerization of the solvent molecules. Energetics (and the rate constant) for this reaction depend strongly on the energetics of ion pairing between the radical anion and the charge compensating cation. When the interaction is strong, as is the case for Li+ ions, the decay of the radical anion is much faster than it is for a large organic cation. Experimentally, this ion pairing proved to be a greater factor controlling the radical anion stability than the effect of substitution in the parent molecule. Following the protonation of the radical anion, multiple reactions occur, but relatively few low molecular weight products have been identified. These products are consistent with the



Page 22 of 42

deprotonation of the solvent and the reductive sulfur extrusion in the anolyte, but there are also polyanion products of electrolyte decomposition with the unknown structure. Our study decreases the likelihood that further lowering of the redox potentials for neutral, closed-shell anolyte molecules is feasible in CH3CN due to the facility of such concerted proton transfer reactions and the strong correlation between the proton affinity and the redox potential in the anolyte molecules. Lowering this redox potential is very likely to shorten the lifetime of the radical anion, and there needs to be a compromise between the two desired performance criteria, the charge carrier stability and the energy density. While the ion pairing of the anolyte radical anions exerts strong effect on their stability, we found no such effect for the catholyte radical cations. The answer to the second question is also complicated: even when one sacrifices the energy density by increasing the redox potential through derivatization with the electron withdrawing groups (in order to slow down the proton transfer) this does not necessarily translate to greater stability of the radical anion, because such groups are more readily cleaved in the radical anion, and such side reactions shorten the radical anion lifetime. Furthermore, as the redox potential becomes more positive, the second redox reaction can occur, making it more difficult to control the cell charging during cycling. 23 All in all, only modest gains in the radical anion stability and the calendar life of the charged fluids were obtained in this way. Thus, our study suggests that there can exist natural limits to improving the stability of neutral, closed-shell anolytes with low redox potentials that cannot be surpassed by tuning their electronic properties through derivatization. While these conclusions have been reached by studying a specific class of anolyte molecules, we believe that they are general, as the reactions involved are also general. We caution that our conclusions pertain to the stability of radical anions in the electrolyte bulk; the electrochemical stability (that also includes reactions occurring near the electrodes and reactions involving products generated during cell cycling) may not follow these stability trends at all due to such compounding factors.

22

We consider these

additional factors in ref. 23


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ASSOCIATED CONTENT Supporting Information: A PDF file containing the list of abbreviations, additional schemes, figures and tables and the synthetic and experimental sections. This material is available free of charge via the Internet at http://pubs.acs.org.

ACKNOWLEDGMENTS This work was supported as part of the Joint Center for Energy Storage Research (JCESR), an Energy Innovation Hub funded by the U.S. Department of Energy, Office of Science, Basic Energy Sciences. The submitted manuscript has been created by UChicago Argonne, LLC, Operator of Argonne National Laboratory (“Argonne”). Argonne, a U.S. Department of Energy Office of Science laboratory, is operated under Contract No. DE-AC02-06CH11357. The U.S. Government retains for itself, and others acting on its behalf, a paid-up nonexclusive, irrevocable worldwide license in said article to reproduce, prepare derivative works, distribute copies to the public, and perform publicly and display publicly, by or on behalf of the Government.



Page 24 of 42

Table 1. BzNSN Derivatives and their Redox Potentials in 0.5 M LiTFSI/ CH3CN.

R4 R5

3 N

S2 R6 R7

N 1

1 2 3 4 5 6 7 8 9

R4

R5

R6

R7

H H H H H H H CN NO2

H Me Me MeO F CF3 MeOCO H H

H H Me H H H H H H

H H H H H H H CN H

a

-E0 vs Ag/Ag+, V 1.590 1.647 1.674 1.681 1.537 1.356 1.370 0.925 b c

a) 5 mM redox active molecules, b) the second redox potential is -1.519 V, c) not reversible.


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Table 2. EPR Parameters for X+1-● Pairs in DME.

X 7

Li Na 39 K NEt4 23

e

Na f

% 92.4 c 100 c 93.3 c 64.5 d 20.0 d 80.0 d

Aat(s), a G 13.0 33.1 8.1 33.1

absolute isotropic hfcc’s, G (1 G = 10-4 T) 14 1 1 N(1,3) H(4,7) H(5,6) 5.453 2.150 1.506 5.349 2.502 1.592 5.278 2.573 1.607 5.144 2.675 1.497 5.276 2.536 1.554 5.320 2.493 1.580

Alk+ 0.214 0.483 0.093 0.492

ρ(s), b x10-2

(g-2), x104

1.65 1.46 1.15 -

55.7 53.5 56.9 64.5 -

a) atomic hfcc for fully occupied outer s-orbital, b) electron density in the outer s-orbital, c) atom% at the natural abundance, d) mol% in solution; e) free 1-● anions in 1:5 v/v DME:THF solution, f) Na+1-● pairs in the same solution.



Page 26 of 42

Table 3. EPR Parameters for Contact 7Li+ BzNSN-● Pairs in 0.5 M LiTFSI/ CH3CN.

R3 H H H H H H CN CN

substitution R4 R5 MeO H Me Me Me H H H F H MeOC(O) H H H H H

R7 H H H H H H CN CN

(g-2) x104 56 61 51 55 54 55 50 50

7

Li(2) 0.193 0.227 0.178 0.205 0.141 0.112 0.044 0.038

absolute isotropic hfcc’s, G (1 G = 10-4 T) 14 N(1) 14N(3) (4) (5) (6) 6.010 4.932 2.194 0.294 1.276 5.521 5.521 1.267 1.661 1.661 5.632 5.314 2.053 1.366 1.792 5.480 5.480 1.937 1.589 1.589 5.858 4.917 2.552 1.648 1.926 3.993 5.628 2.731 0.377 0.134 2.877 2.877 0.807 1.554 1.554 2.786 2.786 0.781 1.591 1.591

(7) 2.204 1.267 2.151 1.937 2.642 2.283 0.807 0.781

Table 4. EPR Parameters for Weakly Bound NBu4+ BzNSN-● Pairs in 0.5 M NBu4PF6/ CH3CN .

R3 H H H

substitution R4 R5 MeO H Me H H H

R7 H H H

(g-2) x104 65 65 65

absolute isotropic hfcc’s, G (1 G = 10-4 T) 14 N(1) 14N(3) (4) (5) (6) (7) 5.894 4.614 2.984 0.265 1.831 2.364 5.501 5.081 2.673 1.595 1.859 2.326 5.271 5.271 2.617 1.543 1.543 2.617


Page 27 of 42 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 captions.

Figure 1. First-derivative cw EPR spectra of chemically reduced 10 mM 1 in DME (microwave power 2 mW, 0.2 G field modulation, 9.434 GHz). These EPR spectra originate from the contact Alk+ 1-● pairs (Alk = 7Li, 23Na, and 39K) in the solution. Figure 2. Space rendering of the optimized geometry (CH3CN)3Li+1-● contact ion pair. Other solvent molecules were treated as polarizable dielectric continuum. Figure 3. First-derivative cw EPR spectra of electrochemically reduced 5 mM 1 (panel a) and 2 (panel b) in CH3CN solutions containing 0.5 M LiTFSI (red traces) and 0.5 M NBu4PF6 (blue traces; 0.2 G field modulation). Pairing with 7Li+ ions in LiTFSI solutions has considerable effect on the hfcc’s in the radical anion which accounts for the observed changes in the EPR patterns. Figure 4. Decay kinetics for electrochemically generated 50 mM 1-● in (a) in solutions containing 0.5 M LiTFSI in CH3CN with and without 50 mM 10, and (b) in electrolyte solutions containing 0.5 M TFSI (in black) and 0.5 M LiPF6 (in red). The times of half-decay for 1●

are given in the plots, and the dashed line in both panels indicates 50% decay. The

smooth lines are exponential fits. Figure 5. Correlation of the logarithm of the radical anion half-life t1/2 for CH3CN-based electrolytes containing 0.5 M salt (indicated in the plot) with the redox potential of BzNSN molecules (filled symbols). Also shown (open circles) are the t1/2 times for 10+●. Figure 6. 27 ACS Paragon Plus Environment


Page 28 of 42

Decay kinetics for electrochemically generated 2-● (a) and 4-● (b) in the cathodic fluids initially containing 50 mM 2 or 4, 50 mM 10, and 0.5 M LiTFSI or NBu4PF6 in CH3CN. The t1/2 values given in the plot indicate the time of half-decay found from the exponential extrapolations (smooth lines).

Figure 7. Decay kinetics for electrochemically generated (a) 5-● and (b) 7-● and 8-● in the cathodic fluids initially containing 50 mM ROM, 50 mM 10, and 0.5 M LiTFSI or NBu4PF6 in CH3CN (as indicated in the plot). For 5 (unlike is the case for 1-4) replacing Li+ by NBu4+ dramatically accelerates the decay of the radical anion. While 7 and 8 both have electron withdrawing groups and high redox potentials, 7-● is short-lived, whereas 8-● is long-lived.


Page 29 of 42

+

-•

Alk 1 in DME

EPR signal, 1st derivative

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


K

Na

Li

3350

3360

3370

3380G

Figure 1. First-derivative cw EPR spectra of chemically reduced 10 mM 1 in DME (microwave power 2 mW, 0.2 G field modulation, 9.434 GHz). These EPR spectra originate from the contact Alk+ 1-● pairs (Alk = 7Li, 23Na, and 39K) in the solution.



Page 30 of 42

Figure 2. Space rendering of the optimized geometry (CH3CN)3Li+1-● contact ion pair. Other solvent molecules were treated as polarizable dielectric continuum.



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



Page 31 of 42

(a) 1

-•

5 mM ROM 0.5 M salt / CH3CN

NBu4PF6

LiTFSI

(b) 2

-•

NBu4PF6

LiTFSI

3340

3350

3360

3370G

Figure 3. First-derivative cw EPR spectra of electrochemically reduced 5 mM 1 (panel a) and 2 (panel b) in CH3CN solutions containing 0.5 M LiTFSI (red traces) and 0.5 M NBu4PF6 (blue traces; 0.2 G field modulation). Pairing with 7Li+ ions in LiTFSI solutions has considerable effect on the hfcc’s in the radical anion which accounts for the observed changes in the EPR patterns.



I(t)/I0

1.0

0.5 M LiTFSI with 10 no 10 t1/2 = 52.2±0.1 h

0.5

1

t1/2 = 51.0±0.1 h

-•

(a) 0.0 1.0

I(t)/I0

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

Page 32 of 42

0.5 M salt LiTFSI LiPF6 t1/2 = 52.2±0.1 h

0.5

1

t1/2 = 45±0.1 h

-•

(b) 0.0 0

3

2

4x10 time, min

Figure 4. Decay kinetics for electrochemically generated 50 mM 1-● in (a) in solutions containing 0.5 M LiTFSI in CH3CN with and without 50 mM 10, and (b) in electrolyte solutions containing 0.5 M TFSI (in black) and 0.5 M LiPF6 (in red). The times of half-decay for 1-● are given in the plots, and the dashed line in both panels indicates 50% decay. The smooth lines are exponential fits.


Page 33 of 42

1000

8 6

radical cation anion

+•

Me 2

NBu4PF6

H

4

t1/2, h

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


10

LiTFSI

MeO

100

8 6

F, 5 H, 1

4

LiPF6 2

Me, 2

10

8 6

2Me, 3 MeO, 4

4

-1.7 -1.6 -1.5 + redox potential vs Ag/Ag

Figure 5. Correlation of the logarithm of the radical anion half-life t1/2 for CH3CN-based electrolytes containing 0.5 M salt (indicated in the plot) with the redox potential of BzNSN molecules (filled symbols). Also shown (open circles) are the t1/2 times for 10+●.



1.0

-•

t1/2 = 319±1 h

I(t)/I0

(a) 2 0.5

t1/2 = 14.8 h 0.0

1.0 0.5 M salt LiTFSI NBu4PF6 I(t)/I0

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

Page 34 of 42

0.5

(b) 4

-•

t1/2 = 184±3 h

t1/2 = 7.4 h 0.0 0

2

4x10

3

time, min

Figure 6. Decay kinetics for electrochemically generated 2-● (a) and 4-● (b) in the cathodic fluids initially containing 50 mM 2 or 4, 50 mM 10, and 0.5 M LiTFSI or NBu4PF6 in CH3CN. The t1/2 values given in the plot indicate the time of half-decay found from the exponential extrapolations (smooth lines).


Page 35 of 42

1.0

(a)

I(t)/I0

5

-•

t1/2 = 92.6±0.3 h 0.5

5

-•

0.5 M salt LiTFSI NBu4PF6

t1/2 = 1.5 h 0.0 0

2

4x10

3

1.0

(b) I(t)/I0

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


7

8

-•

-•

0.5

t1/2 = 5.7 h 0.5 M LiTFSI

0.0 0

5

10x10

3

time, min

Figure 7. Decay kinetics for electrochemically generated (a) 5-● and (b) 7-● and 8-● in the cathodic fluids initially containing 50 mM ROM, 50 mM 10, and 0.5 M LiTFSI or NBu4PF6 in CH3CN (as indicated in the plot). For 5 (unlike is the case for 1-4) replacing Li+ by NBu4+ dramatically accelerates the decay of the radical anion. While 7 and 8 both have electron withdrawing groups and high redox potentials, 7-● is short-lived, whereas 8-● is long-lived.



Page 36 of 42

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Elucidating Factors Controlling Long-Term Stability of Radical Anions

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