Supporting Information Tin Oxide-Dependence of the CO2 Reduction ...
Supporting Information Tin Oxide-Dependence of the CO2 Reduction Efficiency on Tin Electrodes and Enhanced Activity for Tin/Tin Oxide Thin Film Catalysts Yihong Chen and Matthew W. Kanan* *To whom correspondence should be addressed. E-mail: [email protected]
Index
page
Complete reference 3
S2
Experimental methods
S2–S8
Table S1. Total geometric current densities, faradaic efficiencies for CO and HCO2H production on etched Sn foil at selected potentials.
S9
Fig. S1. XPS results of native oxide growth test.
S10
Fig. S2. Survey X-ray photoelectron spectra of untreated Sn foil, etched Sn foil and the Sn/SnOx catalyst on a Ti electrode.
S11– S12
Fig. S3. Powder X-ray diffraction pattern of untreated Sn foil.
S13
Fig. S4. Cyclic voltammetry of Sn/SnOx catalyst in 0.5 M NaHCO3/N2 and 0.5 M NaHCO3/CO2.
S14
Fig. S5. Loss of conductivity in Sn/SnOx after exposure to air.
S15
Fig. S6. Addition of Sn2+ to the electrolyte for an electrolysis with Sn foil
S16
Fig. S7. Electrolysis with Sn foil at –0.7 V before and after electrolysis at –1.4 V
S17
Complete Reference 3 Blankenship, R. E.; Tiede, D. M.; Barber, J.; Brudvig, G. W.; Fleming, G.; Ghirardi, M.; Gunner, M. R.; Junge, W.; Kramer, D. M.; Melis, A.; Moore, T. A.; Moser, C. C.; Nocera, D. G.; Nozik, A. J.; Ort, D. R.; Parson, W. W.; Prince, R. C.; Sayre, R. T. Science 2011, 322, 805. Experimental Methods Materials. SnCl2 (>99.99%), Na2CO3 (>99.9999%), and hydrobromic acid (ACS reagent, 48% in water) were purchased from Sigma Aldrich; carbon dioxide (99.99%) was purchased from Praxair; titanium foil (>99.9%, 0.1 mm thick) was purchased from MTI Corporation; tin foil (99.998%, 0.25 mm thick) was purchased from Alfa Aesar. Unless otherwise stated, all chemicals were used without further purification. Electrolyte solutions were prepared with DI water (Ricca Chemical, ASTM Type I). 0.50 M NaHCO3 solution was prepared by vigorously bubbling CO2 gas through 0.25 M Na2CO3 solution for at least 6 hours. Etching procedures. Tin foil electrodes were etched by soaking in hot dilute HBr solution[1] or by cathodizing in 1:5 HCl:H2O solution at –3.0 V for 5 min as described.[2] Electrochemical measurements. A CH Instruments 760D or 660D potentiostat was used for all experiments. A piece of platinum gauze (Sigma, 99.9%) was used as auxiliary electrode. The electrolyte used for all CO2 reduction experiments was 0.5 M NaHCO3 saturated with CO2. The pH is of this electrolyte is 7.2. All potentials were measured against an Ag/AgCl reference electrode (3.0 M KCl, World Precision Instruments) and converted to the RHE reference scale using E (vs RHE) = E (vs Ag/AgCl) 0.210 V 0.0591 V pH.
CO2 reduction electrolyses and product analysis. Electrolyses were performed in a gas-tight two-compartment electrochemical cell with a piece of SELEMION anion exchange membrane as a separator. All glassware was cleaned in fresh hot Piranha solution prior to experiments. Each compartment contained 20 mL of 0.5 M NaHCO3 electrolyte. The headspace of cathodic compartment was approximately 5 mL. CO2 gas was delivered into the cathodic compartment at a rate of 5.00 sccm and the compartment was vented directly into the gas-sampling loop of a gas chromatograph (GC) (SRI Instruments). A GC run was initiated every 15 minutes. The GC was equipped with a packed MoleSieve 5A column and a packed HayeSep D column. Argon (Praxair, 99.999%) was used as the carrier gas. A flame ionization detector (FID) with methanizer was used to quantity CO concentration and a thermal conductivity detector (TCD) was used to quantify hydrogen concentration. The partial current density of CO production was calculated from the GC peak area as follows:
jCO =
peak area 2Fp0 flow rate (electrode area) –1 RT
where is a conversion factor based on calibration of the GC with a standard sample, p0 = 1.013 bar and T = 273.15 K. Using a CO standard consisting of 991 ppm (v/v) CO in CO2: RT 420 mV• min • sccm = mA 2Fp0
HCO2H concentration was analyzed on a Varian Inova 600 MHz NMR spectrometer. A 0.5 mL sample of the electrolyte was mixed with 0.1 mL D2O and 1.67ppm (m/m) di-
methyl sulfoxide (DMSO, Sigma, 99.99%) added as an internal standard. The 1D 1H spectrum was measured with water suppression using a presaturation method. In situ formation of Sn/SnOx on Ti. A 1.0 cm 2.0 cm piece of Ti foil was pretreated by refluxing in 10% oxalic acid (Sigma, >99%) solution for 60 min. Before electrolysis, the electrolyte in the cathodic compartment was degassed by bubbling with CO2 gas for at least 30 min until the GC analysis showed only CO2 in the exhaust gas. The electrolyte in cathodic compartment was stirred at a rate of 800 rpm during electrolysis. For catalyst deposition, 1 mL of 20 mM SnCl2 or Sn(OTf)2 solution was injected into the cathode chamber shortly after the electrolysis was started. SnCl2 and Sn(OTf)2 solutions were prepared fresh for each experiment. Cyclic voltammetry (CV). Cyclic voltammagrams were obtained in the same cell that was used for electrolyses. CO2 or N2 gas was vigorously bubbled into the solution for at least 30 minutes before the measurements. CV traces of the Sn/SnOx catalyst were obtained after 2 h of bulk electrolysis in NaHCO3/CO2. In N2-saturated 0.5 M NaHCO3 (Figure S4a), Sn foil exhibits a sharp reduction wave at Ep = –0.16 V and a broad oxidation wave beginning at –0.1 V. Based on previous studies in borate electrolytes, the broad oxidation wave is consistent with the formation of multiple Sn2+/4+ oxides and hydroxides that are converted to SnO2 as the potential is increased; the reduction wave corresponds to the Sn4+/0 reduction of SnO2 to Sn0.[3] A similar voltammagram is observed for the Sn/SnOx electrode, however the reduction wave exhibits additional features and a long tail. This difference indicates the presence of a much thicker SnOx layer on this electrode. Similar CV features are observed in CO2-saturated
electrolyte, although the peak positions are significantly shifted (Figure S4b). The comparable total peak areas in the CVs of Sn foil and Sn/SnOx indicate similar total amounts of redox active Sn. Although the CO2 reduction bulk electrolyses with Sn foil and Sn/SnOx are performed at potentials negative of the reduction waves observed in CV, most of the SnOx present on these electrodes is not reduced over the course of a long electrolysis, as evidenced by XPS. This persistence may result from poor electrical contact to the oxide or from electrocatalytic activity of the oxide that effects reduction of H+ and CO2 in solution instead of the oxide itself. Tafel slope. The current density vs. potential data were obtained by stepped-potential electrolyses. The current densities of CO production at different potentials were calculated from the GC spectra every 15 min and averaged. A 0.5 mL aliquot of the electrolyte was extracted at the end of each potential step. Average partial current densities for HCO2H production at each step were calculated from NMR quantification of the HCO2H in these aliquots. Considerable variation of the CO2 reduction efficiency of untreated Sn foil electrodes was observed when several different samples were evaluated. The data used for Sn foil in Figure 3 are for an electrode that was more active than most other samples evaluated. The variability in CO2 reduction activity likely arises from variable amounts of native SnOx on the samples of untreated Sn foil. While initial oxide growth is rapid, maximum oxide growth as determined by XPS requires several weeks of air exposure (see below). Ex situ analyses. Scanning electron microscopy (SEM) images were acquired with FEI Magellan 400 XHR Scanning Electron Microscope. X-ray photoelectron spectra were
obtained with a PHI VersaProbe II Scanning XPS Microprobe. Powder X-ray diffraction (XRD) patterns were obtained with a PANalytical's X'Pert PRO Materials Research Diffractometer with Programmable Divergence Slit (PDS) and PIXcel3D detector. Residual SnCl2 concentration in electrolyte was measured on a Thermo Scientific XSERIES 2 Quadrupole Inductively Coupled Plasma Mass Spectrometer (ICP-MS). Native oxide growth on a Sn0 electrode. A piece of Sn foil was etched in hot HBr solution as described above and subsequently transferred to the XPS chamber. A 2 mm 2 mm section of the electrode was sputtered by treating with Ar+ plasma under ultrahigh vacuum (UHV) for 10 minutes and a high resolution Sn 3d5/2 spectrum was obtained (Figure S1a). The electrode was subsequently extracted from UHV, exposed to air for 10 min in the antechamber of the XPS, and reinserted without changing the sample position. A second high-resolution Sn 3d5/2 X-ray photoelectron spectrum was obtained (Figure S1b). The peak assigned to native oxide (SnOx) increased from 18% to 64% as a result of the 10 min air exposure. Notably, etched electrodes analyzed by XPS after 1 day of exposure to air also exhibited a 60–70% SnOx peak, indicating that the growth of the oxide plateaus. Several weeks of air exposure are required to attain a 90+% SnOx peak. Exposure of Sn/SnOx to air. A Sn/SnOx catalyst was prepared on Ti foil by deposition at –0.7 V in a NaHCO3/CO2 electrolyte containing 1 mM SnCl2. The current density and CO efficiency of the in situ deposited catalyst were typical of Sn/SnOx samples prepared in this manner (Figure S5). After 1 h and 40 min of CO2 reduction electrolysis, the cell was switched to open circuit and the cathode was transferred to a new cell with NaHCO3/CO2 electrolyte. The Sn/SnOx catalyst was exposed to air for 2–3 min for this transfer. Electrolysis was initiated at –0.7 V in the new cell. A very low current density (< 20 μA/cm2)
was observed compared to the current density prior to air exposure. This result is consistent with an oxidation-induced loss of conductivity in the Sn/SnOx film. Since the film is composed of nanoscale particles (Figure 2b), we hypothesize that growth of SnOx layers between the particles in the presence of O2 forms resistive barriers. Addition of Sn2+ to the electrolyte for an electrolysis with untreated Sn foil. An untreated Sn foil electrode with a native SnOx layer was used as the cathode in a CO2 reduction electrolysis at –0.7 V. After 3 h, an aliquot of concentrated SnCl2 in H2O was added to the catholyte such that the final concentration of Sn2+ was 1 mM. The current density rapidly increased from 600 μA/cm2 before addition of Sn2+ to 2 mA/cm2 after addition of Sn2+ and the faradaic efficiency for CO production increased from 15% to 30% (Figure S6). This result is consistent with the deposition of Sn/SnOx on the Sn foil substrate. The faradaic efficiency and partial current density for CO are generally lower for Sn/SnOx formed deposited on Sn compared to Sn/SnOx deposited on Ti and therefore the data in Figures 2 and 3 of the main text are for Sn/SnOx on Ti. We hypothesize that the nucleation of Sn/SnOx is sensitive to the surface of the conductive substrate upon which it is deposited and the background current density of the substrate during the deposition. Alternative substrates may result in more active Sn/SnOx thin film catalysts. Electrolysis with Sn foil at –0.7 V before and after electrolysis at –1.4 V. An untreated Sn foil electrode with a native SnOx layer was used as the cathode in a CO2 reduction electrolysis at –0.7 V. After 2 h 40 min at this potential, the potential was switched to – 1.4 V and electrolysis was allowed to proceed for 20 min. The potential was subsequently switched back to –0.7 V and the electrolysis was continued for an additional 2 h 40 min. The current density and CO faradaic efficiencies at –0.7 V after the electrolysis at –1.4 V
were very similar to the values at –0.7 V before this step (Figure S7). This result indicates that SnOx layers are stable in NaHCO3/CO2 electrolytes under cathodic potentials significantly negative of the potentials required to achieve mass transport–limited CO2 reduction current densities. References [1]
Hsu, Y.-S.; Ghandhi, S. K. J Electrochem Soc 1979, 126, 1434.
[2]
Baliga, B. J.; Ghandhi, S. K. J Electrochem Soc 1977, 124, 1059.
[3]
Díaz, R.; Joiret, S.; Cuesta, Á.; Díez-Pérez, I.; Allongue, P.; Gutiérrez, C.; Gorostiza, P.; Sanz, F. J Phys Chem B 2004, 108, 8173.
Table S1 Table S1. Summary of total geometric current densities, Faradaic efficiencies for CO and HCO2H production on etched Sn foil at different potentials. E vs. RHE / V
Faraday Efficiency / %
jtot / mA cm-2
CO
HCO2H
–1.06
11.9
0.5
Index
page
Complete reference 3
S2
Experimental methods
S2–S8
Table S1. Total geometric current densities, faradaic efficiencies for CO and HCO2H production on etched Sn foil at selected potentials.
S9
Fig. S1. XPS results of native oxide growth test.
S10
Fig. S2. Survey X-ray photoelectron spectra of untreated Sn foil, etched Sn foil and the Sn/SnOx catalyst on a Ti electrode.
S11– S12
Fig. S3. Powder X-ray diffraction pattern of untreated Sn foil.
S13
Fig. S4. Cyclic voltammetry of Sn/SnOx catalyst in 0.5 M NaHCO3/N2 and 0.5 M NaHCO3/CO2.
S14
Fig. S5. Loss of conductivity in Sn/SnOx after exposure to air.
S15
Fig. S6. Addition of Sn2+ to the electrolyte for an electrolysis with Sn foil
S16
Fig. S7. Electrolysis with Sn foil at –0.7 V before and after electrolysis at –1.4 V
S17
Complete Reference 3 Blankenship, R. E.; Tiede, D. M.; Barber, J.; Brudvig, G. W.; Fleming, G.; Ghirardi, M.; Gunner, M. R.; Junge, W.; Kramer, D. M.; Melis, A.; Moore, T. A.; Moser, C. C.; Nocera, D. G.; Nozik, A. J.; Ort, D. R.; Parson, W. W.; Prince, R. C.; Sayre, R. T. Science 2011, 322, 805. Experimental Methods Materials. SnCl2 (>99.99%), Na2CO3 (>99.9999%), and hydrobromic acid (ACS reagent, 48% in water) were purchased from Sigma Aldrich; carbon dioxide (99.99%) was purchased from Praxair; titanium foil (>99.9%, 0.1 mm thick) was purchased from MTI Corporation; tin foil (99.998%, 0.25 mm thick) was purchased from Alfa Aesar. Unless otherwise stated, all chemicals were used without further purification. Electrolyte solutions were prepared with DI water (Ricca Chemical, ASTM Type I). 0.50 M NaHCO3 solution was prepared by vigorously bubbling CO2 gas through 0.25 M Na2CO3 solution for at least 6 hours. Etching procedures. Tin foil electrodes were etched by soaking in hot dilute HBr solution[1] or by cathodizing in 1:5 HCl:H2O solution at –3.0 V for 5 min as described.[2] Electrochemical measurements. A CH Instruments 760D or 660D potentiostat was used for all experiments. A piece of platinum gauze (Sigma, 99.9%) was used as auxiliary electrode. The electrolyte used for all CO2 reduction experiments was 0.5 M NaHCO3 saturated with CO2. The pH is of this electrolyte is 7.2. All potentials were measured against an Ag/AgCl reference electrode (3.0 M KCl, World Precision Instruments) and converted to the RHE reference scale using E (vs RHE) = E (vs Ag/AgCl) 0.210 V 0.0591 V pH.
CO2 reduction electrolyses and product analysis. Electrolyses were performed in a gas-tight two-compartment electrochemical cell with a piece of SELEMION anion exchange membrane as a separator. All glassware was cleaned in fresh hot Piranha solution prior to experiments. Each compartment contained 20 mL of 0.5 M NaHCO3 electrolyte. The headspace of cathodic compartment was approximately 5 mL. CO2 gas was delivered into the cathodic compartment at a rate of 5.00 sccm and the compartment was vented directly into the gas-sampling loop of a gas chromatograph (GC) (SRI Instruments). A GC run was initiated every 15 minutes. The GC was equipped with a packed MoleSieve 5A column and a packed HayeSep D column. Argon (Praxair, 99.999%) was used as the carrier gas. A flame ionization detector (FID) with methanizer was used to quantity CO concentration and a thermal conductivity detector (TCD) was used to quantify hydrogen concentration. The partial current density of CO production was calculated from the GC peak area as follows:
jCO =
peak area 2Fp0 flow rate (electrode area) –1 RT
where is a conversion factor based on calibration of the GC with a standard sample, p0 = 1.013 bar and T = 273.15 K. Using a CO standard consisting of 991 ppm (v/v) CO in CO2: RT 420 mV• min • sccm = mA 2Fp0
HCO2H concentration was analyzed on a Varian Inova 600 MHz NMR spectrometer. A 0.5 mL sample of the electrolyte was mixed with 0.1 mL D2O and 1.67ppm (m/m) di-
methyl sulfoxide (DMSO, Sigma, 99.99%) added as an internal standard. The 1D 1H spectrum was measured with water suppression using a presaturation method. In situ formation of Sn/SnOx on Ti. A 1.0 cm 2.0 cm piece of Ti foil was pretreated by refluxing in 10% oxalic acid (Sigma, >99%) solution for 60 min. Before electrolysis, the electrolyte in the cathodic compartment was degassed by bubbling with CO2 gas for at least 30 min until the GC analysis showed only CO2 in the exhaust gas. The electrolyte in cathodic compartment was stirred at a rate of 800 rpm during electrolysis. For catalyst deposition, 1 mL of 20 mM SnCl2 or Sn(OTf)2 solution was injected into the cathode chamber shortly after the electrolysis was started. SnCl2 and Sn(OTf)2 solutions were prepared fresh for each experiment. Cyclic voltammetry (CV). Cyclic voltammagrams were obtained in the same cell that was used for electrolyses. CO2 or N2 gas was vigorously bubbled into the solution for at least 30 minutes before the measurements. CV traces of the Sn/SnOx catalyst were obtained after 2 h of bulk electrolysis in NaHCO3/CO2. In N2-saturated 0.5 M NaHCO3 (Figure S4a), Sn foil exhibits a sharp reduction wave at Ep = –0.16 V and a broad oxidation wave beginning at –0.1 V. Based on previous studies in borate electrolytes, the broad oxidation wave is consistent with the formation of multiple Sn2+/4+ oxides and hydroxides that are converted to SnO2 as the potential is increased; the reduction wave corresponds to the Sn4+/0 reduction of SnO2 to Sn0.[3] A similar voltammagram is observed for the Sn/SnOx electrode, however the reduction wave exhibits additional features and a long tail. This difference indicates the presence of a much thicker SnOx layer on this electrode. Similar CV features are observed in CO2-saturated
electrolyte, although the peak positions are significantly shifted (Figure S4b). The comparable total peak areas in the CVs of Sn foil and Sn/SnOx indicate similar total amounts of redox active Sn. Although the CO2 reduction bulk electrolyses with Sn foil and Sn/SnOx are performed at potentials negative of the reduction waves observed in CV, most of the SnOx present on these electrodes is not reduced over the course of a long electrolysis, as evidenced by XPS. This persistence may result from poor electrical contact to the oxide or from electrocatalytic activity of the oxide that effects reduction of H+ and CO2 in solution instead of the oxide itself. Tafel slope. The current density vs. potential data were obtained by stepped-potential electrolyses. The current densities of CO production at different potentials were calculated from the GC spectra every 15 min and averaged. A 0.5 mL aliquot of the electrolyte was extracted at the end of each potential step. Average partial current densities for HCO2H production at each step were calculated from NMR quantification of the HCO2H in these aliquots. Considerable variation of the CO2 reduction efficiency of untreated Sn foil electrodes was observed when several different samples were evaluated. The data used for Sn foil in Figure 3 are for an electrode that was more active than most other samples evaluated. The variability in CO2 reduction activity likely arises from variable amounts of native SnOx on the samples of untreated Sn foil. While initial oxide growth is rapid, maximum oxide growth as determined by XPS requires several weeks of air exposure (see below). Ex situ analyses. Scanning electron microscopy (SEM) images were acquired with FEI Magellan 400 XHR Scanning Electron Microscope. X-ray photoelectron spectra were
obtained with a PHI VersaProbe II Scanning XPS Microprobe. Powder X-ray diffraction (XRD) patterns were obtained with a PANalytical's X'Pert PRO Materials Research Diffractometer with Programmable Divergence Slit (PDS) and PIXcel3D detector. Residual SnCl2 concentration in electrolyte was measured on a Thermo Scientific XSERIES 2 Quadrupole Inductively Coupled Plasma Mass Spectrometer (ICP-MS). Native oxide growth on a Sn0 electrode. A piece of Sn foil was etched in hot HBr solution as described above and subsequently transferred to the XPS chamber. A 2 mm 2 mm section of the electrode was sputtered by treating with Ar+ plasma under ultrahigh vacuum (UHV) for 10 minutes and a high resolution Sn 3d5/2 spectrum was obtained (Figure S1a). The electrode was subsequently extracted from UHV, exposed to air for 10 min in the antechamber of the XPS, and reinserted without changing the sample position. A second high-resolution Sn 3d5/2 X-ray photoelectron spectrum was obtained (Figure S1b). The peak assigned to native oxide (SnOx) increased from 18% to 64% as a result of the 10 min air exposure. Notably, etched electrodes analyzed by XPS after 1 day of exposure to air also exhibited a 60–70% SnOx peak, indicating that the growth of the oxide plateaus. Several weeks of air exposure are required to attain a 90+% SnOx peak. Exposure of Sn/SnOx to air. A Sn/SnOx catalyst was prepared on Ti foil by deposition at –0.7 V in a NaHCO3/CO2 electrolyte containing 1 mM SnCl2. The current density and CO efficiency of the in situ deposited catalyst were typical of Sn/SnOx samples prepared in this manner (Figure S5). After 1 h and 40 min of CO2 reduction electrolysis, the cell was switched to open circuit and the cathode was transferred to a new cell with NaHCO3/CO2 electrolyte. The Sn/SnOx catalyst was exposed to air for 2–3 min for this transfer. Electrolysis was initiated at –0.7 V in the new cell. A very low current density (< 20 μA/cm2)
was observed compared to the current density prior to air exposure. This result is consistent with an oxidation-induced loss of conductivity in the Sn/SnOx film. Since the film is composed of nanoscale particles (Figure 2b), we hypothesize that growth of SnOx layers between the particles in the presence of O2 forms resistive barriers. Addition of Sn2+ to the electrolyte for an electrolysis with untreated Sn foil. An untreated Sn foil electrode with a native SnOx layer was used as the cathode in a CO2 reduction electrolysis at –0.7 V. After 3 h, an aliquot of concentrated SnCl2 in H2O was added to the catholyte such that the final concentration of Sn2+ was 1 mM. The current density rapidly increased from 600 μA/cm2 before addition of Sn2+ to 2 mA/cm2 after addition of Sn2+ and the faradaic efficiency for CO production increased from 15% to 30% (Figure S6). This result is consistent with the deposition of Sn/SnOx on the Sn foil substrate. The faradaic efficiency and partial current density for CO are generally lower for Sn/SnOx formed deposited on Sn compared to Sn/SnOx deposited on Ti and therefore the data in Figures 2 and 3 of the main text are for Sn/SnOx on Ti. We hypothesize that the nucleation of Sn/SnOx is sensitive to the surface of the conductive substrate upon which it is deposited and the background current density of the substrate during the deposition. Alternative substrates may result in more active Sn/SnOx thin film catalysts. Electrolysis with Sn foil at –0.7 V before and after electrolysis at –1.4 V. An untreated Sn foil electrode with a native SnOx layer was used as the cathode in a CO2 reduction electrolysis at –0.7 V. After 2 h 40 min at this potential, the potential was switched to – 1.4 V and electrolysis was allowed to proceed for 20 min. The potential was subsequently switched back to –0.7 V and the electrolysis was continued for an additional 2 h 40 min. The current density and CO faradaic efficiencies at –0.7 V after the electrolysis at –1.4 V
were very similar to the values at –0.7 V before this step (Figure S7). This result indicates that SnOx layers are stable in NaHCO3/CO2 electrolytes under cathodic potentials significantly negative of the potentials required to achieve mass transport–limited CO2 reduction current densities. References [1]
Hsu, Y.-S.; Ghandhi, S. K. J Electrochem Soc 1979, 126, 1434.
[2]
Baliga, B. J.; Ghandhi, S. K. J Electrochem Soc 1977, 124, 1059.
[3]
Díaz, R.; Joiret, S.; Cuesta, Á.; Díez-Pérez, I.; Allongue, P.; Gutiérrez, C.; Gorostiza, P.; Sanz, F. J Phys Chem B 2004, 108, 8173.
Table S1 Table S1. Summary of total geometric current densities, Faradaic efficiencies for CO and HCO2H production on etched Sn foil at different potentials. E vs. RHE / V
Faraday Efficiency / %
jtot / mA cm-2
CO
HCO2H
–1.06
11.9
0.5
