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Insights into the electrocatalytic reduction of CO2 on metallic silver surfaces† Toru Hatsukade, Kendra P. Kuhl, Etosha R. Cave, David N. Abram and Thomas F. Jaramillo* The electrochemical reduction of CO2 could allow for a sustainable process by which renewable energy from wind and solar are used directly in the production of fuels and chemicals. In this work we investigated the potential dependent activity and selectivity of the electrochemical reduction of CO2 on metallic silver surfaces under ambient conditions. Our results deepen our understanding of the surface chemistry and provide insight into the factors important to designing better catalysts for the reaction. The high sensitivity of our experimental methods for identifying and quantifying products of reaction
Received 17th February 2014, Accepted 2nd June 2014
allowed for the observation of six reduction products including CO and hydrogen as major products and
DOI: 10.1039/c4cp00692e
behavior of all products, we provide insights into kinetics and mechanisms at play, in particular involving the production of hydrocarbons and alcohols on catalysts with weak CO binding energy as well as the
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formation of a C–C bond required to produce ethanol.
formate, methane, methanol, and ethanol as minor products. By quantifying the potential-dependent
Introduction CO2 emissions from the consumption of fossil fuels for energy have been linked to the anthropogenic climate change, and its mitigation is one of the greatest challenges of our time.1 Replacement of fossil fuels with renewable energy sources such as wind and solar is a promising solution. However, the intermittent nature of these sources is problematic, as the integration of these sources into the grid causes destabilization when they exceed 20–30% of grid capacity.2,3 Electrochemical reduction of CO2 would allow us to overcome these limitations by enabling the storage of renewable energies in the form of fuels and chemicals.4–6 However, one key barrier to its utilization is the lack of efficient and selective catalysts. In order to develop an effective catalyst for the process, fundamental understanding of the catalytic reaction must be established. Previous work in this area has identified several catalyst characteristics that are desirable for the electrochemical reduction of CO2 to hydrocarbons and oxygenates in aqueous media at ambient conditions (25 1C, 1 atm). The first is low activity for the hydrogen evolution reaction (HER), as this reaction tends to compete with the CO2 reduction reaction (CO2RR) due to their close reversible potentials and generally more facile HER kinetics.5 A late onset for the HER would allow
Department of Chemical Engineering, Stanford University, 381 North-South Mall, Stanford, California 94305, USA. E-mail: [email protected] † Electronic supplementary information (ESI) available. See DOI: 10.1039/ c4cp00692e
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for higher selectivity towards the CO2RR in the lower overpotential region, which is vital to achieving high energy efficiency for the CO2RR. The second is optimal CO binding energy (EB[CO]) of the catalyst surface, as CO is commonly proposed as the reactive intermediate for the production of hydrocarbons and oxygenates.5,7 As per the Sabatier principle, one expects that ‘moderate’ binding energy would allow for CO to be hydrogenated without it poisoning the surface. An appropriate value for EB[CO] could result in a pathway towards hydrocarbons and oxygenates with a favorable overpotential.7 Metallic silver surfaces show low activity for the HER and possess an EB[CO] that is fairly weak, resulting in high selectivity of CO2RR for CO.8–12 The late onset for the HER allows for an in-depth analysis of the CO2RR catalytic activity over a wide range of potentials, making it a suitable system for obtaining mechanistic insights, in particular how such a surface could produce hydrocarbons and/or oxygenates and under which conditions such chemistry might be possible. In this work, the catalytic activity of silver surfaces was investigated for the electrochemical reduction of CO2 in aqueous media at ambient conditions. Our experimental setup offers high sensitivity for minor products of CO2RR, as described previously.13 Hydrogen and CO are commonly reported as reduction products when a Ag electrode is used to catalyze CO2RR, and there are several conflicting reports on the observation of formate12 and methane.14 In our studies, in addition to these previously reported products, we observe the potential dependent production of both methanol and ethanol, reported for the first time here on silver surfaces under any operating conditions. By quantifying
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the potential dependence of the partial current density of each product, we aim to shine light onto how silver catalyzes the electrochemical reduction of CO2.
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Experimental All electrochemical experiments were performed in a custom two compartment electrolysis cell previously introduced by Kuhl et al., shown in Fig. 1.13 The cell setup allowed for a large catalyst surface area compared to the electrolyte volume, which gave rise to a high concentration of liquid products in the electrolyte, resulting in greater sensitivity in identifying and quantifying liquid products. Both compartments were filled with 0.1 M KHCO3 electrolyte (Sigma-Aldrich, 99.99% metals basis) and constantly purged with CO2 (5.0, Praxair) at 20 sccm. The CO2 purge was started 30 minutes prior to the experiment to ensure that the electrolyte reached a constant pH. The pH of the electrolyte was measured at 6.8, both before and after the electrochemical experiment. The exit gas stream from the working electrode compartment was connected directly to a gas chromatograph (GC, SRI 8610C in the Multi-Gas #3 configuration) for immediate analysis of the gas phase products. The compartments were separated by an anion exchange membrane (Selemion AMV, AGC Inc.) to mitigate the transport of liquid phase products from the working electrode to the counter electrode where they could be oxidized. A three electrode setup was utilized with Ag/AgCl (Accumet) as the reference electrode and a platinum foil as the counter electrode. Silver foil (Alfa Aesar, 0.1 mm thickness, 99.998% metals basis) was used as the working electrode. The silver surface was mechanically polished (3M, 400 grit) and rinsed with water until no discoloration was observed on the surface prior to every electrochemical experiments. Electrochemical experiments were performed with an electrochemical impedance spectroscopy (EIS) capable channel in a Biologic VMP3 potentiostat. Cell resistance was measured and
Fig. 1
Schematic of the electrochemical cell utilized in this work.
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85% of the ohmic loss was compensated by the potentiostat for all experiments. The remaining 15% of the loss was manually post-corrected in a similar fashion as described in our previous work.13 The data collected vs. Ag/AgCl reference were converted to a reversible hydrogen electrode (RHE) scale; hence all potentials reported here are versus RHE. The quantification of the gas phase products was performed with the gas chromatograph after 5, 23, 41, and 59 minutes during the hour-long electrolysis. The quantification of the liquid phase products was performed after the electrolysis by running 1D 1H NMR (600 MHz, Varian Inova) on the electrolyte sampled from the working electrode compartment. The characterization of the silver working electrode surface was performed before and after an hour-long electrolysis at 1.23 V. X-ray photoelectron spectroscopy (PHI VersaProbe Scanning XPS Microprobe, Al-Ka radiation) and scanning electron microscopy (FEI XL30 Sirion SEM) were utilized to probe for surface impurities and morphological changes, respectively.
Results and discussion One-hour potentiostatic experiments were conducted at 13 different potentials in the range of 0.60 to 1.42 V. The current densities measured at each potential over the hour of electrolysis are plotted as a function of time in Fig. 2a. The figure shows an increase in the magnitude of the cathodic current density with increasing cathodic potential, as expected. At the less negative potentials, i.e. 0.60, 0.67, and 0.75 V, the current density approximately doubles over the course of an hour. However, this increase in the current density only amounts to B100 mA cm 2, which is imperceptible at the more negative potentials where the current densities are orders of magnitude larger. The nature of this small increase in the current density has yet to be determined and it is beyond the scope of this paper. The average current density over the hour is plotted against potential in Fig. 2b, revealing the potential dependence of the overall activity. An exponential increase in current density is observed initially, interrupted by a brief plateau region, followed by another region of increasing current density. A cyclic voltammogram (CV) obtained prior to the potentiostatic experiments is overlaid on the figure for comparison; the potentiostatic data are consistent with the CV. As the brief plateau region (between B 1.1 V and 1.25 V) is observed in the CV as well, it confirms that it is not an artifact of potentiostatic experiments. The likely cause of this plateau in current density will be discussed later. SEM images of the silver working electrode surface before and after an hour-long experiment are shown in Fig. 3a and b. Comparison of the two images shows that the silver surface becomes rougher during the electrolysis experiment. While the image of the pre-run surface shows a smooth surface with some streaks, presumably left by the mechanical polishing pretreatment, the image of the post-run surface shows rougher features with sizes varying from 50–350 nm. XPS spectra of the two surfaces (Fig. 3c) show neither major differences in chemical state nor the presence of impurities, suggesting that the surface
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Fig. 2 (a) Current density vs. time plot for 13 different potentials. The red arrows indicate when the gas phase samples were taken, and the green arrow indicates when the liquid phase sample was taken. (b) Overlay of CV data and average current densities from hour-long potentiostatic experiments.
Fig. 3 SEM images of the silver surface (a) before and (b) after, and (c) XPS spectrum of silver surface ( ) before and ( ) after an hour-long experiment at 1.23 V vs. RHE.
is altered in a purely morphological fashion. This is further corroborated by additional high resolution scans shown in Fig. S1 (ESI†), which show no significant differences between the two surfaces. The reduction of the native silver oxide and its reoxidation are suggested as a possible origin of the surface roughening. However, further investigation is necessary to determine the true origin of the morphological changes. In the CO2 electrolysis experiments, a total of six reduction products were observed. The major products were H2 and CO, as expected, and interestingly four different minor products were observed: formate, methane, methanol, and ethanol. The current efficiency of each reduction product is plotted against the applied potential in Fig. 4, and the precision of the data is shown in Fig. S2 (ESI†). The plot shows that hydrogen and CO have a combined current efficiency above 90% over the entire potential range. However, the selectivity among these two major
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Fig. 4 Faradaic efficiency for each product as a function of potential.
products changes significantly as a function of potential; hydrogen is dominant at the lowest overpotentials, CO overtakes hydrogen in the intermediate overpotential region, and hydrogen regains dominance at the highest overpotentials. In contrast, the current efficiencies of the minor products (o10% current efficiency) simply increase with overpotential. Formate reaches a maximum current efficiency of 7%, while methane, methanol, and ethanol show current efficiencies ranging from 0.01 to 0.1%; these three particular products are only observed at the highest overpotentials. The observation of methane, methanol, and ethanol at such small amounts and only at the highest overpotentials is likely due to the many proton and electron transfer steps required for each of these products (8, 6, and 12 electrons, respectively), compared to just 2 protons and electrons required for CO and formate. The similar onset potentials of the reduction products requiring 42 electrons suggest the possibility of a common rate determining step.
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Overall, the current efficiencies of the products add up to nearly 100% at each potential as shown in Fig. S3 in the ESI,† which suggests the adequacy of our experimental setup and detection methods. The potential dependent production of methanol and ethanol is notable because it has not been reported previously on silver surfaces under any operating conditions, an observation that we attribute to the high sensitivity of our experimental system for liquid products. The observation of ethanol is particularly intriguing because two-carbon products are rarely observed on metal surfaces at STP in aqueous electrolyte, other than on copper.15 Additional reduction experiments were performed in order to identify the origin of these products, in which we used 13C-labeled CO2 for electrolysis. The 1H NMR spectrum of products from these experiments (Fig. S4, ESI†) shows the 13C splitting of the product peaks, confirming that the products originated from CO2. This information is important to exclude the possibility of these minor products originating from non-CO2 carbon sources (e.g. contamination). The partial current densities of the reduction products are obtained by multiplying the total current density at each potential by the current efficiency of each product, and they are presented as Tafel plots in Fig. 5. This presentation of the data is useful in obtaining insights into the kinetics and mechanisms of the reduction reactions because partial current density is directly proportional to turnover frequency (TOF). To facilitate discussion, we will describe observations within three distinct potential regions of interest: low overpotential ( 0.6 to 1.0 V vs. RHE), intermediate overpotential ( 1.0 to 1.2 V vs. RHE), and high overpotential ( 1.2 to 1.4 V vs. RHE). It is important to note that the designation of the ‘‘low’’ overpotential region is relative, as 40.5 V of overpotential is applied in this region vs. the calculated standard reduction potentials of CO2RR.16 In the low overpotential region, only three reduction products are observed: hydrogen, CO, and formate, each of which requires only 2 electrons to be produced. The TOF of CO increases as a function of overpotential with a slope of ca. 140 mV per decade, whereas the TOF of hydrogen remains constant in this potential region.
Fig. 5 Tafel plot of the partial current density going to each product. Indications for the low, intermediate, and high overpotential regions are shown above the plot.
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This reflects the observed shift in selectivity from hydrogen to CO in this potential region, as discussed earlier in Fig. 4. The TOF of formate also increases as a function of overpotential, but not with a strong dependence. Its fairly flat Tafel slope (ca. 275 mV per decade) and small partial current density at 0.65 V (ca. 2 mA cm 2) prevents it from overtaking the rate of hydrogen production. In the intermediate overpotential region, the TOF of CO suddenly reaches a plateau, which causes a brief plateau to also appear in the total current density plot, since CO is the dominant product within this potential region. Also worthy to note in this potential region is that the H2 TOF remains constant and the formate TOF continues to rise with increasing overpotential, but the formate TOF still remains low compared to H2 or CO. In the high overpotential region, a decrease in CO TOF and a sudden increase in H2 TOF are observed, which explain the shift in selectivity back to hydrogen in this potential region. The products that require more than two electrons (42 e products) begin to emerge within this region; their TOFs increase steadily with overpotential. The change in the Tafel slope of CO observed in the intermediate overpotential region is unexpected and hints at the existence of new factors that come into play. To explore the possibility of mass transfer limitations, the rate of CO2 reduction (molCO2 s 1cm 2) is presented in Fig. 6. The reaction rate vs. potential profile is initially exponential, consistent with the observed Tafel behavior (i.e. a kinetically controlled region), but then shows a dramatic deviation and flattens out, a characteristic shape observed for electrocatalysts switching from kinetic control to mass transport control at higher overpotentials. The existence of mass transport limitations at higher overpotentials was corroborated through additional experimentation (ESI,† Fig. S5a and b), motivating future work in developing electrochemical reactors with greater mass transport properties to extract kinetic values at high overpotential. There remain many important, open questions regarding the mechanism of CO2 electro-reduction on transition metal surfaces, with recent theoretical work shedding new light onto
Fig. 6
Total rate of CO2 reduction as a function of potential.
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Fig. 7 Possible mechanistic pathways for electrochemical reduction of CO2 on silver. (H+ + e ) signifies a reduction step either by distinct proton and electron transfers or by a surface adsorbed H atom.
the matter.7,17,18 To facilitate a discussion based on our experimentally measured current–voltage profiles and associated product selectivities from this work, Fig. 7 illustrates several plausible mechanistic pathways of CO2 electro-reduction on silver including reaction steps involving proton and electron transfers, denoted (H+ + e ), as well as steps involving adsorption, desorption, and C–C coupling that may be either chemical or electrochemical in nature. The first steps (rxns 1a and 1b) depicted in Fig. 7 involve the initial proton and electron transfers to CO2 to form an adsorbed species, either *OCHO (rxn 1a) or *COOH (rxn 1b), where protonation can occur at different atoms of the CO2 molecule, either at the carbon atom or at one of the oxygen atoms, respectively. Either pathway/reaction intermediate could lead to formate by either a one-electron transfer or by means of a proton and electron transfer followed by the loss of a proton to the solution (rxn 2a or 2b). To produce CO, theory suggests rxn 2c whereby a proton and electron transfer to the hydroxyl group in *COOH produces water (H2O), leaving *CO adsorbed to the surface (rxn 2c).7 As *CO is weakly bound to Ag surfaces,19,20 it is expected to desorb readily and emerge as a major reaction product. As discussed earlier, the observation that several 42 e products emerge at similar overpotentials may be due to similarities in their mechanistic pathways, and perhaps indicate a common rate-determining step (r.d.s.). Carbon monoxide is commonly hypothesized as the key intermediate in the pathway towards 42 e products, and its reduction (rxn 3) might be the r.d.s. considering silver’s weak binding of CO. Such weak binding implies rapid desorption once *CO is generated on the surface, leaving little opportunity for further reduction to occur.7 Thus high overpotentials would be necessary to expedite proton and
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electron transfers to generate reaction products with 42 e , since electron transfer steps from the electrode are strongly dependent on applied potential, unlike CO desorption, which is a purely chemical step. The initial observation of a 42 e product (methane) at large overpotentials, 1.23 V, indicates that the rate of CO reduction is beginning to compete with CO desorption, though CO desorption remains dominant even at these very cathodic potentials. In principle, increasing the overpotential even further could eventually result in CO reduction being favored over desorption, but such potentials are well outside the range investigated in this work. In the reduction of *CO, the *CHO intermediate was calculated to be more thermodynamically stable on silver,7 though on copper the *COH intermediate was calculated to be more kinetically accessible.17 To reflect the two possible pathways, both intermediates are shown in Fig. 7. Ethanol formation on silver is particularly intriguing as there are very few surfaces that have been shown to be capable of forming C–C bonds in the electrochemical reduction of CO2.5 Based on the results above, one can infer that the formation of the C–C bond might be a facile reaction step on silver surfaces, given the low surface coverage of C1 intermediates one might expect from the fast desorption kinetics of CO. Mechanistically, two major pathways have been proposed for C–C bond formation in the CO2RR: one involves an electrochemical dimerization of adsorbed *CO, and the other a thermochemical coupling of two adsorbed C1 species. The formation of a *CO dimer mediated by an electron transfer has been suggested as the rate determining step in the electrochemical pathway on copper.21 On the other hand, the reduction of *CO has been suggested as the rate determining step in the thermochemical coupling pathway on copper, as calculations show that the thermochemical coupling of C1 surface intermediates can only occur after CO reduction.22 In both pathways, the rate determining step is electrochemical in nature, consistent with the observed potential dependence of ethanol production measured in our work. However, the fact that ethanol, methane, and methanol share similar onset potentials on silver is more consistent with expectations of the thermochemical pathway; otherwise ethanol would have a different r.d.s. and thus likely a different onset potential. The data are thus consistent with the three products all sharing the same rate determining step, the electrochemical reduction of an absorbed *CO, illustrated as reaction 3 in Fig. 7. From this point, an additional three proton and electron transfers can lead to methanol (reaction 4), while five proton and electron transfers can produce methane (reaction 6). To produce ethanol (reaction 5), a second *C1 species is needed for the C–C coupling process, along with an appropriate number of proton and electron transfers to reach a total of twelve for the molecule as a whole. While further investigation will certainly be necessary to understand the nature of C–C bond formation and the selectivity among the C2 products, the identification of silver as a surface that is capable of performing C–C bond formation chemistry is particularly important to note, especially at ambient conditions.
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While the high overpotential region is interesting due to the production of 42 e CO2RR products, hydrogen production is still dominant in this region. One reason for the high selectivity towards HER in this region is the transport limitations of CO2. However, even if it is possible to increase CO2 transport through reactor engineering and/or operating conditions, competition from the HER is difficult to avoid as proton reduction also becomes more favorable at higher overpotentials. Given this, as well as the low energetic efficiency of processes that require high overpotentials, developing new catalyst materials selective for C–C coupled products at low overpotentials remains an important challenge in the field.
Conclusions We have reported the potential dependence of the catalytic activity of silver surfaces under CO2 reduction conditions. In addition to products commonly reported previously in literature, methanol and ethanol were observed as reduction products. The resulting data were utilized to gain insights into the reaction mechanism, including the overpotentials required for the production of 42 e products, the pathway for the formation of C–C bond in ethanol, the competition between CO2RR and HER, and mass transport effects. While the results suggest that silver is not the most active catalyst for the electrochemical reduction of CO2, its ability to produce a C–C coupled product is particularly intriguing. Further studies will be aimed to elucidate why and how silver is capable of catalyzing the production of this important product.
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Acknowledgements
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This material is based upon work supported by the National Science Foundation under Grant Number 1066515 and by the Global Climate & Energy Project (GCEP) at Stanford University. KPK and ERC acknowledge support by the National Science Foundation Graduate Research Fellowship under Grant No. (DGE-1147470). ERC acknowledges support from the Ford Foundation. DNA acknowledges support from a Stanford Graduate Fellowship. The authors would also like to thank Joey Montoya, Dr Heine A. Hansen, Dr Jakob Kibsgaard, and Prof. Jens K. Norskov for helpful discussions. Thanks to Dr Stephen R. Lynch of the Stanford Chemistry department NMR facility for help with the 600 MHz and 500 MHz experiments.
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References 1 M. I. Hoffert, K. Caldeira, G. Benford, D. R. Criswell, C. Green, H. Herzog, A. K. Jain, H. S. Kheshgi, K. S. Lackner, J. S. Lewis,
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H. D. Lightfoot, W. Manheimer, J. C. Mankins, M. E. Mauel, L. J. Perkins, M. E. Schlesinger, T. Volk and T. M. L. Wigley, Science, 2002, 298, 981–987. H. Chen, T. N. Cong, W. Yang, C. Tan, Y. Li and Y. Ding, Prog. Nat. Sci., 2009, 19, 291–312. N. S. Lewis and D. G. Nocera, Proc. Natl. Acad. Sci. U. S. A., 2006, 103, 15729–15735. M. Gattrell, N. Gupta and A. Co, Energy Convers. Manage., 2007, 48, 1255–1265. Y. Hori, Handbook of Fuel Cells, John Wiley & Sons, Ltd, 2010. D. T. Whipple and P. J. A. Kenis, J. Phys. Chem. Lett., 2010, 1, 3451–3458. A. A. Peterson and J. K. Nørskov, J. Phys. Chem. Lett., 2012, 3, 251–258. K. Watanabe, U. Nagashima and H. Hosoya, Chem. Phys. Lett., 1993, 209, 109–110. K. Watanabe, U. Nagashima and H. Hosoya, Appl. Surf. Sci., 1994, 75, 121–124. Y. Hori, in Modern Aspects of Electrochemistry, ed. C. Vayenas, R. White and M. Gamboa-Aldeco, Springer, New York, 2008, vol. 42, pp. 89–189. J. Greeley, T. F. Jaramillo, J. Bonde, I. B. Chorkendorff and J. K. Nørskov, Nat. Mater., 2006, 5, 909–913. N. Hoshi, M. Kato and Y. Hori, J. Electroanal. Chem., 1997, 440, 283–286. K. P. Kuhl, E. R. Cave, D. N. Abram and T. F. Jaramillo, Energy Environ. Sci., 2012, 5, 7050–7059. H. Noda, S. Ikeda, Y. Oda, K. Imai, M. Maeda and K. Ito, Bull. Chem. Soc. Jpn., 1990, 63, 2459–2462. Y. Hori, H. Wakebe, T. Tsukamoto and O. Koga, Electrochim. Acta, 1994, 39, 1833–1839. in CRC Handbook of Chemistry and Physics, ed. W. M. Haynes, CRC Press/Taylor and Francis, Boca Raton, FL, 93rd edn (internet version 2013), 2013. X. Nie, M. R. Esopi, M. J. Janik and A. Asthagiri, Angew. Chem., Int. Ed., 2013, 52, 2459–2462. F. Calle-Vallejo and M. T. Koper, Angew. Chem., 2013, 52, 7282–7285. H. Falsig, B. Hvolbæk, I. S. Kristensen, T. Jiang, T. Bligaard, C. H. Christensen and J. K. Nørskov, Angew. Chem., Int. Ed., 2008, 47, 4835–4839. T. Jiang, D. J. Mowbray, S. Dobrin, H. Falsig, B. Hvolbæk, T. Bligaard and J. K. Nørskov, J. Phys. Chem. C, 2009, 113, 10548–10553. K. J. P. Schouten, Y. Kwon, C. J. M. van der Ham, Z. Qin and M. T. M. Koper, Chem. Sci., 2011, 2, 1902. J. H. Montoya, A. A. Peterson and J. K. Nørskov, ChemCatChem, 2013, 5, 737–742.
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Insights into the electrocatalytic reduction of CO2 on metallic silver surfaces† Toru Hatsukade, Kendra P. Kuhl, Etosha R. Cave, David N. Abram and Thomas F. Jaramillo* The electrochemical reduction of CO2 could allow for a sustainable process by which renewable energy from wind and solar are used directly in the production of fuels and chemicals. In this work we investigated the potential dependent activity and selectivity of the electrochemical reduction of CO2 on metallic silver surfaces under ambient conditions. Our results deepen our understanding of the surface chemistry and provide insight into the factors important to designing better catalysts for the reaction. The high sensitivity of our experimental methods for identifying and quantifying products of reaction
Received 17th February 2014, Accepted 2nd June 2014
allowed for the observation of six reduction products including CO and hydrogen as major products and
DOI: 10.1039/c4cp00692e
behavior of all products, we provide insights into kinetics and mechanisms at play, in particular involving the production of hydrocarbons and alcohols on catalysts with weak CO binding energy as well as the
www.rsc.org/pccp
formation of a C–C bond required to produce ethanol.
formate, methane, methanol, and ethanol as minor products. By quantifying the potential-dependent
Introduction CO2 emissions from the consumption of fossil fuels for energy have been linked to the anthropogenic climate change, and its mitigation is one of the greatest challenges of our time.1 Replacement of fossil fuels with renewable energy sources such as wind and solar is a promising solution. However, the intermittent nature of these sources is problematic, as the integration of these sources into the grid causes destabilization when they exceed 20–30% of grid capacity.2,3 Electrochemical reduction of CO2 would allow us to overcome these limitations by enabling the storage of renewable energies in the form of fuels and chemicals.4–6 However, one key barrier to its utilization is the lack of efficient and selective catalysts. In order to develop an effective catalyst for the process, fundamental understanding of the catalytic reaction must be established. Previous work in this area has identified several catalyst characteristics that are desirable for the electrochemical reduction of CO2 to hydrocarbons and oxygenates in aqueous media at ambient conditions (25 1C, 1 atm). The first is low activity for the hydrogen evolution reaction (HER), as this reaction tends to compete with the CO2 reduction reaction (CO2RR) due to their close reversible potentials and generally more facile HER kinetics.5 A late onset for the HER would allow
Department of Chemical Engineering, Stanford University, 381 North-South Mall, Stanford, California 94305, USA. E-mail: [email protected] † Electronic supplementary information (ESI) available. See DOI: 10.1039/ c4cp00692e
13814 | Phys. Chem. Chem. Phys., 2014, 16, 13814--13819
for higher selectivity towards the CO2RR in the lower overpotential region, which is vital to achieving high energy efficiency for the CO2RR. The second is optimal CO binding energy (EB[CO]) of the catalyst surface, as CO is commonly proposed as the reactive intermediate for the production of hydrocarbons and oxygenates.5,7 As per the Sabatier principle, one expects that ‘moderate’ binding energy would allow for CO to be hydrogenated without it poisoning the surface. An appropriate value for EB[CO] could result in a pathway towards hydrocarbons and oxygenates with a favorable overpotential.7 Metallic silver surfaces show low activity for the HER and possess an EB[CO] that is fairly weak, resulting in high selectivity of CO2RR for CO.8–12 The late onset for the HER allows for an in-depth analysis of the CO2RR catalytic activity over a wide range of potentials, making it a suitable system for obtaining mechanistic insights, in particular how such a surface could produce hydrocarbons and/or oxygenates and under which conditions such chemistry might be possible. In this work, the catalytic activity of silver surfaces was investigated for the electrochemical reduction of CO2 in aqueous media at ambient conditions. Our experimental setup offers high sensitivity for minor products of CO2RR, as described previously.13 Hydrogen and CO are commonly reported as reduction products when a Ag electrode is used to catalyze CO2RR, and there are several conflicting reports on the observation of formate12 and methane.14 In our studies, in addition to these previously reported products, we observe the potential dependent production of both methanol and ethanol, reported for the first time here on silver surfaces under any operating conditions. By quantifying
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the potential dependence of the partial current density of each product, we aim to shine light onto how silver catalyzes the electrochemical reduction of CO2.
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Experimental All electrochemical experiments were performed in a custom two compartment electrolysis cell previously introduced by Kuhl et al., shown in Fig. 1.13 The cell setup allowed for a large catalyst surface area compared to the electrolyte volume, which gave rise to a high concentration of liquid products in the electrolyte, resulting in greater sensitivity in identifying and quantifying liquid products. Both compartments were filled with 0.1 M KHCO3 electrolyte (Sigma-Aldrich, 99.99% metals basis) and constantly purged with CO2 (5.0, Praxair) at 20 sccm. The CO2 purge was started 30 minutes prior to the experiment to ensure that the electrolyte reached a constant pH. The pH of the electrolyte was measured at 6.8, both before and after the electrochemical experiment. The exit gas stream from the working electrode compartment was connected directly to a gas chromatograph (GC, SRI 8610C in the Multi-Gas #3 configuration) for immediate analysis of the gas phase products. The compartments were separated by an anion exchange membrane (Selemion AMV, AGC Inc.) to mitigate the transport of liquid phase products from the working electrode to the counter electrode where they could be oxidized. A three electrode setup was utilized with Ag/AgCl (Accumet) as the reference electrode and a platinum foil as the counter electrode. Silver foil (Alfa Aesar, 0.1 mm thickness, 99.998% metals basis) was used as the working electrode. The silver surface was mechanically polished (3M, 400 grit) and rinsed with water until no discoloration was observed on the surface prior to every electrochemical experiments. Electrochemical experiments were performed with an electrochemical impedance spectroscopy (EIS) capable channel in a Biologic VMP3 potentiostat. Cell resistance was measured and
Fig. 1
Schematic of the electrochemical cell utilized in this work.
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85% of the ohmic loss was compensated by the potentiostat for all experiments. The remaining 15% of the loss was manually post-corrected in a similar fashion as described in our previous work.13 The data collected vs. Ag/AgCl reference were converted to a reversible hydrogen electrode (RHE) scale; hence all potentials reported here are versus RHE. The quantification of the gas phase products was performed with the gas chromatograph after 5, 23, 41, and 59 minutes during the hour-long electrolysis. The quantification of the liquid phase products was performed after the electrolysis by running 1D 1H NMR (600 MHz, Varian Inova) on the electrolyte sampled from the working electrode compartment. The characterization of the silver working electrode surface was performed before and after an hour-long electrolysis at 1.23 V. X-ray photoelectron spectroscopy (PHI VersaProbe Scanning XPS Microprobe, Al-Ka radiation) and scanning electron microscopy (FEI XL30 Sirion SEM) were utilized to probe for surface impurities and morphological changes, respectively.
Results and discussion One-hour potentiostatic experiments were conducted at 13 different potentials in the range of 0.60 to 1.42 V. The current densities measured at each potential over the hour of electrolysis are plotted as a function of time in Fig. 2a. The figure shows an increase in the magnitude of the cathodic current density with increasing cathodic potential, as expected. At the less negative potentials, i.e. 0.60, 0.67, and 0.75 V, the current density approximately doubles over the course of an hour. However, this increase in the current density only amounts to B100 mA cm 2, which is imperceptible at the more negative potentials where the current densities are orders of magnitude larger. The nature of this small increase in the current density has yet to be determined and it is beyond the scope of this paper. The average current density over the hour is plotted against potential in Fig. 2b, revealing the potential dependence of the overall activity. An exponential increase in current density is observed initially, interrupted by a brief plateau region, followed by another region of increasing current density. A cyclic voltammogram (CV) obtained prior to the potentiostatic experiments is overlaid on the figure for comparison; the potentiostatic data are consistent with the CV. As the brief plateau region (between B 1.1 V and 1.25 V) is observed in the CV as well, it confirms that it is not an artifact of potentiostatic experiments. The likely cause of this plateau in current density will be discussed later. SEM images of the silver working electrode surface before and after an hour-long experiment are shown in Fig. 3a and b. Comparison of the two images shows that the silver surface becomes rougher during the electrolysis experiment. While the image of the pre-run surface shows a smooth surface with some streaks, presumably left by the mechanical polishing pretreatment, the image of the post-run surface shows rougher features with sizes varying from 50–350 nm. XPS spectra of the two surfaces (Fig. 3c) show neither major differences in chemical state nor the presence of impurities, suggesting that the surface
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Fig. 2 (a) Current density vs. time plot for 13 different potentials. The red arrows indicate when the gas phase samples were taken, and the green arrow indicates when the liquid phase sample was taken. (b) Overlay of CV data and average current densities from hour-long potentiostatic experiments.
Fig. 3 SEM images of the silver surface (a) before and (b) after, and (c) XPS spectrum of silver surface ( ) before and ( ) after an hour-long experiment at 1.23 V vs. RHE.
is altered in a purely morphological fashion. This is further corroborated by additional high resolution scans shown in Fig. S1 (ESI†), which show no significant differences between the two surfaces. The reduction of the native silver oxide and its reoxidation are suggested as a possible origin of the surface roughening. However, further investigation is necessary to determine the true origin of the morphological changes. In the CO2 electrolysis experiments, a total of six reduction products were observed. The major products were H2 and CO, as expected, and interestingly four different minor products were observed: formate, methane, methanol, and ethanol. The current efficiency of each reduction product is plotted against the applied potential in Fig. 4, and the precision of the data is shown in Fig. S2 (ESI†). The plot shows that hydrogen and CO have a combined current efficiency above 90% over the entire potential range. However, the selectivity among these two major
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Fig. 4 Faradaic efficiency for each product as a function of potential.
products changes significantly as a function of potential; hydrogen is dominant at the lowest overpotentials, CO overtakes hydrogen in the intermediate overpotential region, and hydrogen regains dominance at the highest overpotentials. In contrast, the current efficiencies of the minor products (o10% current efficiency) simply increase with overpotential. Formate reaches a maximum current efficiency of 7%, while methane, methanol, and ethanol show current efficiencies ranging from 0.01 to 0.1%; these three particular products are only observed at the highest overpotentials. The observation of methane, methanol, and ethanol at such small amounts and only at the highest overpotentials is likely due to the many proton and electron transfer steps required for each of these products (8, 6, and 12 electrons, respectively), compared to just 2 protons and electrons required for CO and formate. The similar onset potentials of the reduction products requiring 42 electrons suggest the possibility of a common rate determining step.
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Overall, the current efficiencies of the products add up to nearly 100% at each potential as shown in Fig. S3 in the ESI,† which suggests the adequacy of our experimental setup and detection methods. The potential dependent production of methanol and ethanol is notable because it has not been reported previously on silver surfaces under any operating conditions, an observation that we attribute to the high sensitivity of our experimental system for liquid products. The observation of ethanol is particularly intriguing because two-carbon products are rarely observed on metal surfaces at STP in aqueous electrolyte, other than on copper.15 Additional reduction experiments were performed in order to identify the origin of these products, in which we used 13C-labeled CO2 for electrolysis. The 1H NMR spectrum of products from these experiments (Fig. S4, ESI†) shows the 13C splitting of the product peaks, confirming that the products originated from CO2. This information is important to exclude the possibility of these minor products originating from non-CO2 carbon sources (e.g. contamination). The partial current densities of the reduction products are obtained by multiplying the total current density at each potential by the current efficiency of each product, and they are presented as Tafel plots in Fig. 5. This presentation of the data is useful in obtaining insights into the kinetics and mechanisms of the reduction reactions because partial current density is directly proportional to turnover frequency (TOF). To facilitate discussion, we will describe observations within three distinct potential regions of interest: low overpotential ( 0.6 to 1.0 V vs. RHE), intermediate overpotential ( 1.0 to 1.2 V vs. RHE), and high overpotential ( 1.2 to 1.4 V vs. RHE). It is important to note that the designation of the ‘‘low’’ overpotential region is relative, as 40.5 V of overpotential is applied in this region vs. the calculated standard reduction potentials of CO2RR.16 In the low overpotential region, only three reduction products are observed: hydrogen, CO, and formate, each of which requires only 2 electrons to be produced. The TOF of CO increases as a function of overpotential with a slope of ca. 140 mV per decade, whereas the TOF of hydrogen remains constant in this potential region.
Fig. 5 Tafel plot of the partial current density going to each product. Indications for the low, intermediate, and high overpotential regions are shown above the plot.
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This reflects the observed shift in selectivity from hydrogen to CO in this potential region, as discussed earlier in Fig. 4. The TOF of formate also increases as a function of overpotential, but not with a strong dependence. Its fairly flat Tafel slope (ca. 275 mV per decade) and small partial current density at 0.65 V (ca. 2 mA cm 2) prevents it from overtaking the rate of hydrogen production. In the intermediate overpotential region, the TOF of CO suddenly reaches a plateau, which causes a brief plateau to also appear in the total current density plot, since CO is the dominant product within this potential region. Also worthy to note in this potential region is that the H2 TOF remains constant and the formate TOF continues to rise with increasing overpotential, but the formate TOF still remains low compared to H2 or CO. In the high overpotential region, a decrease in CO TOF and a sudden increase in H2 TOF are observed, which explain the shift in selectivity back to hydrogen in this potential region. The products that require more than two electrons (42 e products) begin to emerge within this region; their TOFs increase steadily with overpotential. The change in the Tafel slope of CO observed in the intermediate overpotential region is unexpected and hints at the existence of new factors that come into play. To explore the possibility of mass transfer limitations, the rate of CO2 reduction (molCO2 s 1cm 2) is presented in Fig. 6. The reaction rate vs. potential profile is initially exponential, consistent with the observed Tafel behavior (i.e. a kinetically controlled region), but then shows a dramatic deviation and flattens out, a characteristic shape observed for electrocatalysts switching from kinetic control to mass transport control at higher overpotentials. The existence of mass transport limitations at higher overpotentials was corroborated through additional experimentation (ESI,† Fig. S5a and b), motivating future work in developing electrochemical reactors with greater mass transport properties to extract kinetic values at high overpotential. There remain many important, open questions regarding the mechanism of CO2 electro-reduction on transition metal surfaces, with recent theoretical work shedding new light onto
Fig. 6
Total rate of CO2 reduction as a function of potential.
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Fig. 7 Possible mechanistic pathways for electrochemical reduction of CO2 on silver. (H+ + e ) signifies a reduction step either by distinct proton and electron transfers or by a surface adsorbed H atom.
the matter.7,17,18 To facilitate a discussion based on our experimentally measured current–voltage profiles and associated product selectivities from this work, Fig. 7 illustrates several plausible mechanistic pathways of CO2 electro-reduction on silver including reaction steps involving proton and electron transfers, denoted (H+ + e ), as well as steps involving adsorption, desorption, and C–C coupling that may be either chemical or electrochemical in nature. The first steps (rxns 1a and 1b) depicted in Fig. 7 involve the initial proton and electron transfers to CO2 to form an adsorbed species, either *OCHO (rxn 1a) or *COOH (rxn 1b), where protonation can occur at different atoms of the CO2 molecule, either at the carbon atom or at one of the oxygen atoms, respectively. Either pathway/reaction intermediate could lead to formate by either a one-electron transfer or by means of a proton and electron transfer followed by the loss of a proton to the solution (rxn 2a or 2b). To produce CO, theory suggests rxn 2c whereby a proton and electron transfer to the hydroxyl group in *COOH produces water (H2O), leaving *CO adsorbed to the surface (rxn 2c).7 As *CO is weakly bound to Ag surfaces,19,20 it is expected to desorb readily and emerge as a major reaction product. As discussed earlier, the observation that several 42 e products emerge at similar overpotentials may be due to similarities in their mechanistic pathways, and perhaps indicate a common rate-determining step (r.d.s.). Carbon monoxide is commonly hypothesized as the key intermediate in the pathway towards 42 e products, and its reduction (rxn 3) might be the r.d.s. considering silver’s weak binding of CO. Such weak binding implies rapid desorption once *CO is generated on the surface, leaving little opportunity for further reduction to occur.7 Thus high overpotentials would be necessary to expedite proton and
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electron transfers to generate reaction products with 42 e , since electron transfer steps from the electrode are strongly dependent on applied potential, unlike CO desorption, which is a purely chemical step. The initial observation of a 42 e product (methane) at large overpotentials, 1.23 V, indicates that the rate of CO reduction is beginning to compete with CO desorption, though CO desorption remains dominant even at these very cathodic potentials. In principle, increasing the overpotential even further could eventually result in CO reduction being favored over desorption, but such potentials are well outside the range investigated in this work. In the reduction of *CO, the *CHO intermediate was calculated to be more thermodynamically stable on silver,7 though on copper the *COH intermediate was calculated to be more kinetically accessible.17 To reflect the two possible pathways, both intermediates are shown in Fig. 7. Ethanol formation on silver is particularly intriguing as there are very few surfaces that have been shown to be capable of forming C–C bonds in the electrochemical reduction of CO2.5 Based on the results above, one can infer that the formation of the C–C bond might be a facile reaction step on silver surfaces, given the low surface coverage of C1 intermediates one might expect from the fast desorption kinetics of CO. Mechanistically, two major pathways have been proposed for C–C bond formation in the CO2RR: one involves an electrochemical dimerization of adsorbed *CO, and the other a thermochemical coupling of two adsorbed C1 species. The formation of a *CO dimer mediated by an electron transfer has been suggested as the rate determining step in the electrochemical pathway on copper.21 On the other hand, the reduction of *CO has been suggested as the rate determining step in the thermochemical coupling pathway on copper, as calculations show that the thermochemical coupling of C1 surface intermediates can only occur after CO reduction.22 In both pathways, the rate determining step is electrochemical in nature, consistent with the observed potential dependence of ethanol production measured in our work. However, the fact that ethanol, methane, and methanol share similar onset potentials on silver is more consistent with expectations of the thermochemical pathway; otherwise ethanol would have a different r.d.s. and thus likely a different onset potential. The data are thus consistent with the three products all sharing the same rate determining step, the electrochemical reduction of an absorbed *CO, illustrated as reaction 3 in Fig. 7. From this point, an additional three proton and electron transfers can lead to methanol (reaction 4), while five proton and electron transfers can produce methane (reaction 6). To produce ethanol (reaction 5), a second *C1 species is needed for the C–C coupling process, along with an appropriate number of proton and electron transfers to reach a total of twelve for the molecule as a whole. While further investigation will certainly be necessary to understand the nature of C–C bond formation and the selectivity among the C2 products, the identification of silver as a surface that is capable of performing C–C bond formation chemistry is particularly important to note, especially at ambient conditions.
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While the high overpotential region is interesting due to the production of 42 e CO2RR products, hydrogen production is still dominant in this region. One reason for the high selectivity towards HER in this region is the transport limitations of CO2. However, even if it is possible to increase CO2 transport through reactor engineering and/or operating conditions, competition from the HER is difficult to avoid as proton reduction also becomes more favorable at higher overpotentials. Given this, as well as the low energetic efficiency of processes that require high overpotentials, developing new catalyst materials selective for C–C coupled products at low overpotentials remains an important challenge in the field.
Conclusions We have reported the potential dependence of the catalytic activity of silver surfaces under CO2 reduction conditions. In addition to products commonly reported previously in literature, methanol and ethanol were observed as reduction products. The resulting data were utilized to gain insights into the reaction mechanism, including the overpotentials required for the production of 42 e products, the pathway for the formation of C–C bond in ethanol, the competition between CO2RR and HER, and mass transport effects. While the results suggest that silver is not the most active catalyst for the electrochemical reduction of CO2, its ability to produce a C–C coupled product is particularly intriguing. Further studies will be aimed to elucidate why and how silver is capable of catalyzing the production of this important product.
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Acknowledgements
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This material is based upon work supported by the National Science Foundation under Grant Number 1066515 and by the Global Climate & Energy Project (GCEP) at Stanford University. KPK and ERC acknowledge support by the National Science Foundation Graduate Research Fellowship under Grant No. (DGE-1147470). ERC acknowledges support from the Ford Foundation. DNA acknowledges support from a Stanford Graduate Fellowship. The authors would also like to thank Joey Montoya, Dr Heine A. Hansen, Dr Jakob Kibsgaard, and Prof. Jens K. Norskov for helpful discussions. Thanks to Dr Stephen R. Lynch of the Stanford Chemistry department NMR facility for help with the 600 MHz and 500 MHz experiments.
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