Efficient Electrocatalytic reduction of carbon dioxide in 1-ethyl-3

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Environmental and Carbon Dioxide Issues

Efficient Electrocatalytic reduction of carbon dioxide in 1-ethyl-3methylimidazolium trifluoromethanesulfonate and water mixtures Amanuel Hailu, and Scott K Shaw Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.8b02750 • Publication Date (Web): 19 Oct 2018 Downloaded from http://pubs.acs.org on October 25, 2018

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TITLE: Efficient electrocatalytic reduction of carbon dioxide in 1-ethyl-3-methylimidazolium trifluoromethanesulfonate and water mixtures AUTHORS: Amanuel Hailu Scott K. Shaw

[email protected] [email protected]

INSTITUTION: Department of Chemistry, University of Iowa, Iowa City, Iowa 52242 ABSTRACT: The catalytic activity of room temperature ionic liquid 1-ethyl-3-methylimidazolium trifluoromethanesulfonate ([EMIM][OTf]) for carbon dioxide (CO2) reduction is reported on a polycrystalline silver electrode as functions of water content and solution pH. The optimal aqueous dilution was determined to be ca. 10% v/v, resulting in a reduction of the overpotential required to reduce CO2 in [EMIM][OTf] by 400 mV. The reduction process and products are largely insensitive to pH. CO2 reduction in [EMIM][OTf] with 10% 0.1 M NaHCO3 is achieved with high selectivity for CO with 93.0 ± 4.6 % Faradaic efficiency. Hydrogen evolution reaction (HER) is suppressed in the range of neat [EMIM][OTf] (500 ppm water content) to [EMIM][OTf]:H2O mixture of 10 % water content. Conductivity and viscosity of the [EMIM][OTf]:H2O mixtures suggest that the ionic liquid ion pair fully dissociates; analogous to dilute KCl solution, where each ion is completely hydrated by water molecules. INTRODUCTION: The concentration of atmospheric carbon dioxide (CO2) has increased steadily since the start of the industrial era. CO2 produced by consuming fossil fuels is currently the most significant source of greenhouse gas. Approximately 7 gigatons of carbon dioxide is released each year by fossil fuel burning and deforestation.1-4 To simultaneously address the increasing need for high energy density fuels and offset atmospheric carbon dioxide concentrations, the reduction of carbon dioxide to useful fuels is an attractive idea. From this vantage point, CO2 is a naturally abundant, inexpensive, non-flammable, and non-toxic C1 feedstock. Electrochemistry provides a convenient method to carry out the reduction of carbon dioxide due to its high achievable conversion efficiency. 5-6 Over the past few decades many catalysts have been employed to reduce CO2 electrochemically. Hori et al. reported the selective reduction of CO2 to carbon monoxide (CO) on polycrystalline gold electrode with Faradaic efficiency of 87% at -0.7 V vs reversible hydrogen electrode (RHE).7 Savéant et al. reported the selective reduction of CO2 to CO with Faradaic yields over 90% at -1.16 V vs RHE employing a modified iron tetraphenylporphyrin catalyst.8 Polyansky et al. reported the selective reduction of CO2 to CO with Faradaic yield of 95% at -0.6 V vs RHE over a silver nanostructured electrode.9


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CO2 is thermodynamically stable and kinetically inert. The main hurdle in the electrochemical CO2 reduction lies in the first step, the one-electron reduction of CO2 to form an anion radical (CO2 −•). This activation of CO2 to form CO2 −• requires a high reduction potential of −1.9 V vs NHE, due to a large reorganizational energy between the linear CO2 molecule and, conceivably, the bent CO2 −• anion. Stabilizing the radical anion intermediate formed during CO2 reduction can conceivably lower the overpotential required for this reaction.6 Such stabilization has been proposed for heterogeneous catalysts10 and for a select few solvents, i.e. ionic liquids.11 Carbon dioxide electro-catalysis in room temperature ionic liquids (RTILs) has attracted considerable attention. Imidazolium-based ionic liquids that can strongly interact with CO2 stand out in this regard.11-16 Due to their low vapor pressure, non-flammability, solvating ability, and chemical stability; RTILs have been regarded as “green” alternatives for conventional organic solvents in many chemical processes.17 CO2 is more soluble in RTILs than in water or organic solvents, and RTIL’s natural ionic conductivity and wide potential windows makes them an attractive media for electrochemical reduction of CO2.18 The use of RTILs, especially those featuring cations that interact with CO2, may be advantageous in terms of (i) lowering the overpotential for CO2 reduction by activating CO2 via complexation19, and (ii) suppressing the hydrogen evolution reaction (HER) on the surface of the electrode by site blocking.20-22 Carbon dioxide reduction product distributions and their respective yields strongly depend on the electrode material. Silver is a promising electro-catalyst due to its capability of reducing carbon dioxide to carbon monoxide (CO) selectively.23-25 The selective conversion of CO2 to CO is a promising route for clean energy, and the CO product can be used directly in the Fischer–Tropsch process to produce synthetic fuels from syngas.26 The traditional polycrystalline silver electro-catalyst, however, requires a large overpotential.27-28 To address this limitation we report efforts to carry out the reduction of CO2 in mixtures of imidazolium-based RTIL in water (H2O) as functions IL:H2O solution ratio and solution pH. RTIL electrolytes are expensive relative to water or organic solvents, and many are hygroscopic. Development of CO2 reduction methods based on electrolytes tolerant of significant amounts of water is desirable for large-scale processes. Optimization of solution conditions can offer reduced CO2 reduction overpotentials, hydrogen evolution reaction (HER) suppression, increases in dissolved carbon dioxide diffusion rates, and overall improved current efficiencies. In this work we systematically vary two key parameters of solution conditions, the concentration and source of protons, over three concentration regimes. We discuss our results in the context of earlier results obtained by carrying out the CO2 reduction in aqueous/IL solutions at various pH values. EXPERIMENTAL SECTION Salts and Solvents: 1-Ethyl-3-methylimidazolium trifluoromethanesulfonate ([Emim][OTf]) of purity ≥ 99% (IoLiTec, Tuscaloosa, AL). Trifluoromethanesulfonic acid (HOTf) ≥ 99%, hydrochloric acid (HCl), potassium chloride (KCl), sodium bicarbonate (NaHCO3) (SigmaAldrich, St. Louis, MO) were used as received. Aqueous solutions were prepared with Milli-Q water (18.2 MΩ cm−1). Instruments and Procedures: All experiments are carried at room temperature and ambient pressure. Prior to use, the ionic liquids are dried under vacuum to remove traces of water and stored


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under ultra-pure dinitrogen (Praxair Inc., Danbury, CT) or in a nitrogen glove box with 7, Equation (1) can be written as: 𝐶𝑂2 + 𝑂𝐻 ― ↔𝐻𝐶𝑂3― ,𝐾3 = 10 ―71/𝑀 And finally, the autoprotolysis of water: 𝐻2𝑂↔𝐻 + + 𝑂𝐻 ― , 𝐾𝑊 = 10 ―14𝑀2 Conductivity and viscosity measurements were also carried out for all water dilution solutions in the range 0.05% to 100% (Figure 4 and see Supporting Information, SI-4). The conductivity of 0.05% water content was measured to be close to the literature value of 8.87 ± .008 mS/cm for the neat IL. The plot shows the maximum conductivity of 43.9 ± 0.08 mS/cm at 50 % water dilution.


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This corresponds to the IL concentration of 2.7 mol/L or 10.5 mole ratio (H2O:IL) (Figure 4). Conductivity is proportional to the mobility of ions in the solution and we attribute the increase observed here to the reduction in solution viscosity and to the decrease in the electrostatic interactions between anion and cation. Assuming, as the ion pair interaction is weakened and each ion is hydrated with water molecules, the respective ion size decreases and thus increasing their mobility. The IL solutions’ viscosity depends strongly on their concentration as shown in (see Supporting Information, SI-4). At 0.05 % the viscosity of [EMIM][OTf] was found to be close to the literature value of 41.56 ± 1.42 cP. However, as the water content is increased the viscosity decreases rapidly and then more gradually until the viscosity approaches the viscosity of pure water ~1 cP. From the conductivity and viscosity data a log-log plot of the molar conductivity vs the inverse viscosity is plotted to yield a Walden plot (Figure 5). The plot is compared to the 'ideal' case of 0.01 M KCl which exists as purely dissociated ions. The plot of the [EMIM][OTf] lies parallel but slightly offset to the aqueous KCl line, suggesting that the ionic liquid is composed of mostly independent mobile ions.48-50 The relationship between molar conductivity (Λ) and viscosity (η) can be explained by Walden’s rule, derived from the Stokes–Einstein equation, which connects the diffusion coefficient (D) to the inverse viscosity (η) of the solution medium, and from the Nernst–Einstein equation, which relates the diffusion coefficient to the equivalent conductivity. Here 𝑘 is Boltzmann’s constant, T is temperature in Kelvin (K), 𝑅 is the gas constant, 𝑟 is the hydrodynamic radius, 𝑧 is valence of an ion, 𝐹 is the drag force, 𝑁𝐴 is Avogadros number, and 𝑒0 is the electronic charge. The resulting empirical rule states that the product of molar conductivity and viscosity is a constant at a given temperature.51 kT 𝑅𝑇 D= = 2Λ 6𝜋𝑟η 𝑧𝐹 𝑧𝐹2kT 1 Λ= 6𝜋𝑅𝑇 𝑟η Since 𝐹 = 𝑁𝐴𝑒0 and 𝑘/𝑅 = 1/ 𝑁𝐴, then Λ=

Λη =

𝑧𝑒0F 1 6𝜋𝑇 𝑟η

𝑧𝑒0F 6𝜋𝑟𝑇

= 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡

The molar conductivity value depends on the dilution of an electrolyte. Kohlrausch’s equation can be utilized to determine the limiting molar conductivity or the molar conductivity at infinite dilution. Here Λ0 is the molar conductivity at infinite dilution, A is a constant, and C is the electrolyte concentration.51 Λ = Λ0 ―AC1/2


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A plot of the molar conductivity of [EMIM][OTf] was plotted against the square root of the IL concentration (see Supporting Information, SI-6). The Λ0 was determined by extrapolation to infinite dilution, the extrapolated value was found to be 54.35 ± 0.51 S*cm2*mol-1. Similar results have been reported for other binary IL:H20 mixtures for 1-ethyl-3-methylimidazolium ethylsulfate and 1-ethyl-3-methylimidazolium 2-(2-methoxyethoxy) ethylsulfate.52 Effects of pH: In the interest of obtaining hydrogenated reduced products, effects of proton concentrations (pH) are also investigated. Trifluoromethanesulfonic acid (HOTf) diluted to a concentration of 0.001 M (pH 3) and 0.1 M sodium bicarbonate (NaHCO3) (pH 8) aqueous solutions are used as co-solvents to create aqueous mixtures with [EMIM]OTf] at desired pH values across this range. Similar to results obtained by earlier dilutions with pure water; the optimal IL:water mixture is determined to be 10 % v/v for both pH conditions (see Supporting Information, SI-7 and SI-8). These results suggest that the reduction of CO2 in our system is largely insensitive to proton concentration. Due to slightly improved current densities obtained with 0.1M NaHCO3 as a co-solvent (Figure 6), 10 % dilution of [EMIM][OTf] with 0.1 M NaHCO3 is selected as the optimal condition to reduce CO2. The choice of an appropriate electrolyte media is a major factor to consider in CO2RR. In comparing 0.01 M trifluoromethanesulfonic acid, 0.1 M sodium bicarbonate, and pure water; the bicarbonate solution offers a few additional benefits. NaHCO3, being a buffered electrolyte, can maintain the electrode surface pH close to the bulk value during the electrolysis. Maintaining the bulk pH and the electrode surface pH can help minimize polarization loss. Polarization loss arises due to transport of species (by migration and diffusion) and concentration gradients; the loss is generally the summation of ohmic, diffusion, and Nernstian loss. The ohmic loss is due to the resistance of the solution and the diffusion loss is attributed to the ionic gradient near the electrode surface due to the applied potential.46, 53 CO2 electrolysis: Bulk electrolysis of [EMIM][OTf] with 10 % 0.1 M NaHCO3 under CO2 saturated is carried out in an H-type cell at -1.6 V vs Ag/AgCl for 10 minutes. After this time a 10 µL sample of headspace gas was removed and analyzed with gas chromatography (GC). GC analysis vs pure standards reveal carbon monoxide (CO) as the sole reduced product with an impressive 93.0 ± 4.6 % Faradaic efficiency, similar to other values reported for CO2 reduction in RTILs in the literature.36, 37 The concentration of 0.1 M aqueous NaHCO3 co-solvent was increased from 10 to 100 v/v % at 10 % increments. Bulk electrolysis was carried out at each 10 % increment (Figure 7 and Table 1). At 20% 0.1 M NaHCO3 dilution GC analysis revealed both CO and H2 as reduced products, this supports our previous assignment of the reduced current density observed at 20 % v/v dilution due to competition with HER. CONCLUSION The catalytic activity of [EMIM][OTf] for CO2 reduction is explored as a function of water content and solution pH. The overpotential required to reduce CO2 in [EMIM][OTf] is reduced by ca. 400 mV upon 10 % v/v aqueous dilution. The reduction process and products are largely insensitive to pH concentration. The reduction of CO2 in [EMIM][OTf] with 10 % 0.1 M NaHCO3 is achieved with high selectivity for CO with 93.0 ± 4.6 % Faradaic efficiency. Linear sweep voltammetry scans show that the current density decreases past 10 v/v % dilution, reduced product analysis


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show that past 10 v/v % dilution, H2 gas is a reduced product along with CO, suggesting that the hydrogen evolution reaction starts competing with CO2 reduction past a certain aqueous dilution. A Walden plot of the conductivity and viscosity of varies [EMIM][OTf]:H2O mixtures suggest that the ionic liquid dissociates completely; analogous to dilute KCl solution where each ion is completely hydrated by water molecules. The disassociation of the ionic liquid ion pair is disadvantageous for CO2 reduction as the IL is no longer able to suppress the competing HER. Supporting Information: Supporting information includes: Linear sweep voltammograms (LSV) of CO2 reduction in [EMIM][OTf] with 500-9600 ppm water content (Figure S1), LSV of CO2 reduction in [EMIM][OTf]:H2O mixtures 1-15 % (v/v) (Figure S2), LSV of CO2 reduction in [EMIM][OTf]:H2O mixtures 10-100 % (v/v) (Figure S3), Viscosity measurements of [EMIM][OTf] with increasing percentage of water co-solvent (Figure S4), calculated diffusion coefficient values for CO2 in [EMIM][OTf] with increasing mol fraction of water co-solvent (Figure S5), molar conductivity vs. square root of [EMM][OTf] concentration for IL:H2O mixtures (Figure S6), LSV of CO2 reduction in [EMIM][OTf]:1 mM triflic acid (pH 3) mixtures 10-100 % (v/v) (Figure S7), LSV of CO2 reduction in [EMIM][OTf]:0.1 M NaHCO3 (pH 8) mixtures 10-100 % (v/v) (Figure S8), Faradaic efficiency calculations for bulk electrolysis, calibration curve for CO in [EMIM][OTf]with 10 % v/v of 0.1 M NaHCO3 (pH 8) co-solvent (Figure S9), and calibration curve for H2 in [EMIM][OTf]with 10 % v/v of 0.1 M NaHCO3 (pH 8) co-solvent (Figure S10). Acknowledgements: The authors gratefully acknowledge funding from the American Chemical Society’s Petroleum Research Fund (PRF# 55279-DNI5). We would like to thank Prof. Lou Messerle at the University of Iowa for loaning a gas chromatograph used for the reduction product analysis, and Mr. Benjamin Revis for fabrication of the custom electrochemical cell used here. References Cited: 1. Le Quere, C.; Andrew, R. M.; Canadell, J. G.; Sitch, S.; Korsbakken, J. I.; Peters, G. P.; Manning, A. C.; Boden, T. A.; Tans, P. P.; Houghton, R. A.; Keeling, R. F.; Alin, S.; Andrews, O. D.; Anthoni, P.; Barbero, L.; Bopp, L.; Chevallier, F.; Chini, L. P.; Ciais, P.; Currie, K.; Delire, C.; Doney, S. C.; Friedlingstein, P.; Gkritzalis, T.; Harris, I.; Hauck, J.; Haverd, V.; Hoppema, M.; Goldewijk, K. K.; Jain, A. K.; Kato, E.; Kortzinger, A.; Landschutzer, P.; Lefevre, N.; Lenton, A.; Lienert, S.; Lombardozzi, D.; Melton, J. R.; Metzl, N.; Millero, F.; Monteiro, P. M. S.; Munro, D. R.; Nabel, J. E. M. S.; Nakaoka, S.; O'Brien, K.; Olsen, A.; Omar, A. M.; Ono, T.; Pierrot, D.; Poulter, B.; Rodenbeck, C.; Salisbury, J.; Schuster, U.; Schwinger, J.; Seferian, R.; Skjelvan, I.; Stocker, B. D.; Sutton, A. J.; Takahashi, T.; Tian, H. Q.; Tilbrook, B.; van der Laan-Luijkx, I. T.; van der Werf, G. R.; Viovy, N.; Walker, A. P.; Wiltshire, A. J.; Zaehle, S., Global Carbon Budget 2016. Earth Syst Sci Data 2016, 8 (2), 605-649. 2. Tian, H. Q.; Lu, C. Q.; Ciais, P.; Michalak, A. M.; Canadell, J. G.; Saikawa, E.; Huntzinger, D. N.; Gurney, K. R.; Sitch, S.; Zhang, B. W.; Yang, J.; Bousquet, P.; Bruhwiler, L.; Chen, G. S.;


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36. Fumino, K.; Reimann, S.; Ludwig, R., Probing molecular interaction in ionic liquids by low frequency spectroscopy: Coulomb energy, hydrogen bonding and dispersion forces. Phys Chem Chem Phys 2014, 16 (40), 21903-21929. 37. Koddermann, T.; Fumino, K.; Ludwig, R.; Lopes, J. N. C.; Padua, A. A. H., What FarInfrared Spectra Can Contribute to the Development of Force Fields for Ionic Liquids Used in Molecular Dynamics Simulations. Chemphyschem 2009, 10 (8), 1181-1186. 38. Dong, K.; Zhang, S. J.; Wang, J. J., Understanding the hydrogen bonds in ionic liquids and their roles in properties and reactions. Chem Commun 2016, 52 (41), 6744-6764. 39. Cadogan, S. P.; Maitland, G. C.; Trusler, J. P. M., Diffusion Coefficients of CO2 and N-2 in Water at Temperatures between 298.15 K and 423.15 K at Pressures up to 45 MPa. J Chem Eng Data 2014, 59 (2), 519-525. 40. Morgan, D.; Ferguson, L.; Scovazzo, P., Diffusivities of gases in room-temperature ionic liquids: Data and correlations obtained using a lag-time technique. Ind Eng Chem Res 2005, 44 (13), 4815-4823. 41. Bard, A. J. F., L. R., Electrochemical methods: fundamentals and applications. John Wiley & Sons: Hoboken, 2007. 42. Wasserscheid, P. W., T, Ionic Liquids in Synthesis, Second Edition. Wiley-VCH: Weinheim, 2008. 43. Delacourt, C.; Ridgway, P. L.; Newman, J., Mathematical Modeling of CO2 Reduction to CO in Aqueous Electrolytes I. Kinetic Study on Planar Silver and Gold Electrodes. J Electrochem Soc 2010, 157 (12), B1902-B1910. 44. Delacourt, C.; Newman, J., Mathematical Modeling of CO2 Reduction to CO in Aqueous Electrolytes II. Study of an Electrolysis Cell Making Syngas (CO + H-2) from CO2 and H2O Reduction at Room Temperature. J Electrochem Soc 2010, 157 (12), B1911-B1926. 45. Gupta, N.; Gattrell, M.; MacDougall, B., Calculation for the cathode surface concentrations in the electrochemical reduction of CO2 in KHCO3 solutions. J Appl Electrochem 2006, 36 (2), 161-172. 46. Hashiba, H. W., L.-C.; Chen, Y.; Sato, H. K.; Yotsuhashi, S.; Xiang, C.; Weber, A. Z., Effects of Electrolyte Buffer Capacity on Surface Reactant Species and the Reaction Rate of CO2 in Electrochemical CO2 Reduction. The Journal of Physical Chemistry C 2018, 122 (7), 37193726. 47. Soli, A. L.; Byrne, R. H., CO2 system hydration and dehydration kinetics and the equilibrium CO2/H2CO3 ratio in aqueous NaCl solution. Mar Chem 2002, 78 (2-3), 65-73. 48. MacFarlane, D. R.; Forsyth, M.; Izgorodina, E. I.; Abbott, A. P.; Annat, G.; Fraser, K., On the concept of ionicity in ionic liquids. Phys Chem Chem Phys 2009, 11 (25), 4962-4967. 49. Schreiner, C.; Zugmann, S.; Hartl, R.; Gores, H. J., Fractional Walden Rule for Ionic Liquids: Examples from Recent Measurements and a Critique of the So-Called Ideal KCl Line for the Walden Plot. J Chem Eng Data 2010, 55 (5), 1784-1788. 50. Ries, L. A. S.; do Amaral, F. A.; Matos, K.; Martini, E. M. A.; de Souza, M. O.; de Souza, R. F., Evidence of change in the molecular organization of 1-n-butyl-3-methylimidazolium tetrafluoroborate ionic liquid solutions with the addition of water. Polyhedron 2008, 27 (15), 32873293. 51. Bockris, J. O. M. R., A. K. N. , Modern electrochemistry: an introduction to an interdisciplinary area. Plenum: New York, 1973.


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52. Wong, C. L.; Soriano, A. N.; Li, M. H., Diffusion coefficients and molar conductivities in aqueous solutions of 1-ethyl-3-methylimidazolium-based ionic liquids. Fluid Phase Equilibr 2008, 271 (1-2), 43-52. 53. Singh, M. R.; Clark, E. L.; Bell, A. T., Effects of electrolyte, catalyst, and membrane composition and operating conditions on the performance of solar-driven electrochemical reduction of carbon dioxide. Phys Chem Chem Phys 2015, 17 (29), 18924-18936.


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

Current Density (mA/cm2)

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10 0 -10 -20 -30 -2.5 -2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0 Potential (V vs Ag/AgCl)

Figure 1: Electrochemical window of [EMIM][OTf] (500 ppm water) on polycrystalline Ag working electrode (2 mm diameter) between −2.4 and +0.8 V vs Ag/AgCl under inert atmosphere (N2 purge), five consecutive cycles. Scan rate: 50 mV/s.


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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

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Current density (mA/cm2)

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N2

0 -5

CO2

-10 -15 -20 -25

-2.5

-2.0 -1.5 -1.0 -0.5 Potential (V vs Ag/AgCl)

0.0

Figure 2: LSVs obtained of [EMIM][OTf] on a Ag working electrode (2 mm diameter) in saturated N2 (black) and CO2 (red), 500 ppm water content. Scan rate: 50mV/s.


-1.6 -1.7 -1.8

Peak potential (V) Viscosity (cP)

-1.9 -2.0 -2.1

0

2

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44 40 36 32 28 24 20 16 12 8

Viscosity (cP)

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

Average peak potential (V vs Ag/AgCl)

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4 6 8 10 12 14 16 Water content % (v/v)

Figure 3: Plot of carbon dioxide reduction peak potential as a function of [EMIM][OTf] with increasing percentage of water co-solvent ranging from 0.05-15 % (v/v) in saturated CO2 (standard deviation as error bars for n>3).


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:DWHU FRQWHQW

YY

&RQGXFWLYLW\ P6 FP

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

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0ROH UDWLR + 2 >(0,0@>27I@

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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

Log(Molar conductivity (S*cm2 mol-1))

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Page 18 of 21

3.0 2.5 2.0

EMIM OTf 0.01M KClaq Linear fit

KClaq

1.5 1.0 0.5 0.0

y=(-0.132)+(0.918)*x

-0.5

R2=0.998

-1.0 -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 3.0 Log(1/Viscosity) (Poise-1)

Figure 5: A log-log plot of molar conductivity vs. viscosity indicates the degree of ionicity of the ionic liquid. Plots are compared to the 'ideal' case of 0.01 M KCl which exists as dissociated ions. Data collected at 25 ºC.


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0 Current Density (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

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-2 -4 -6

10 % (v/v) H2O 10 % (v/v) 0.1 M NaHCO3 10 % (v/v) 1 mM HOTf

-8 -2.0

-1.6 -1.2 -0.8 -0.4 Potential (V vs Ag/AgCl)

0.0

Figure 6: Overlay of LSVs obtained of [EMIM][OTf] on a Ag working electrode (2 mm diameter) water in saturated CO2 with 10 v/v% milliQ water (black), 10 v/v% 0.1 M sodium bicarbonate (pH 8) (red), and 10 v/v% 1 mM triflic acid (pH 3) (blue) co-solvent. Scan rate: 50 mV/s.


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

Faradaic Efficiency (%)

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100 90 80 70 60 50 40 30 20 10 0 0

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CO H2 20

40 60 80 100 NaHCO3 co-solvent (v/v %)

Figure 7: Faradaic efficiencies for CO and H2 attained with controlled control electrolysis carried out at -1.6 V vs Ag/AgCl for 10 minutes in [EMIM][OTf]with increasing v/v% of 0.1 M NaHCO3 (pH 8) co-solvent (error bars are σ for n>3).


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Table 1: Faradaic efficiencies for CO and H2 attained with controlled electrolysis carried out at -1.6 V vs Ag/AgCl for 10 minutes in [EMIM][OTf]with increasing v/v% of 0.1 M NaHCO3 (pH 8) co-solvent. 0.1M NaHCO3 co-solvent (v/v%)

Carbon monoxide (CO) FE%

Hydrogen (H2) FE%

10

93 ± 5

N/A

20

83 ± 6

1±1

30

49 ± 2

19 ± 1

40

40 ± 3

36 ± 4

50

31 ± 2

56 ± 2

60

32 ± 3

65 ± 3

70

34 ± 4

65 ± 3

80

29 ± 1

65 ± 2

90

31 ± 2

76 ± 3

100

24 ± 5

84 ± 3


Efficient Electrocatalytic reduction of carbon dioxide in 1-ethyl-3

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