Alkaline CO2 Electrolysis toward Selective and Continuous HCOO

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Alkaline CO Electrolysis Towards Selective and Continuous HCOO Production over SnO Nanocatalysts -

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Seunghwa Lee, Joey Duran Ocon, Young-il Son, and Jaeyoung Lee J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp512436w • Publication Date (Web): 10 Feb 2015 Downloaded from http://pubs.acs.org on February 18, 2015

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The Journal of Physical Chemistry

Alkaline CO2 Electrolysis towards Selective and Continuous HCOOProduction over SnO2 Nanocatalysts Seunghwa Lee,† Joey D. Ocon,† Young-il Son,§ Jaeyoung Lee*,†,‡ †

Electrochemical Reaction and Technology Laboratory (ERTL), School of Environmental Science and Technology, ‡Ertl Center for Electrochemistry and Catalysis, RISE, Gwangju Institute of Science and Technology (GIST), Gwangju 500-712, South Korea § Korea Environmental Industry & Technology, Seoul 122-706, South Korea.

ABSTRACT

The electrolyte pH is an important parameter in determining the equilibrium concentrations of the carbon dioxide-bicarbonate-carbonate system, as well as in mapping out the thermodynamically stable phases of tin dioxide (SnO2) in an aqueous electrochemical system. Thus, we explored an optimized region in the combined potential-pH (E-pH) diagram of the two systems, where there is a simultaneously high catalytic activity for carbon dioxide (CO2) electrolysis and good phase stability for the SnO2 nanocatalysts. Our results suggest that choosing the right E-pH combination, which in this case is at 0.6 V (vs. RHE) and pH=10.2, resulted in a high faradaic efficiency of 67.6 % for formate (HCOO-) synthesis and an efficiency retention of ~90% after 5 hr, while maintaining the stability of the oxide structure and avoiding the formation of carbon monoxide. Widely applicable to neutral or near-neutral pH metal oxide electrocatalysts, optimized alkaline CO2 electrolysis offer distinct advantages in terms of the three major catalyst properties: activity, selectivity, and stability.

KEYWORDS: CO2 reduction, SnO2 catalyst, HCOO-/HCOOH, Alkaline electrolysis, Stability

ACS Paragon Plus Environment

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1. Introduction In recent decades, unabated carbon dioxide (CO2) emissions from fossil fuel combustion and other anthropogenic activities have captured great attention due to its adverse effects in the world’s climate. In view of reducing CO2 in the atmosphere, the use of CO2 as a widely available C1 feedstock for the synthesis of valuable substances is put forward as more beneficial over CO2 capture and sequestration techniques.

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Among various conversion methods, the

electrochemical route to synthesizing high value chemicals from CO2 offers several advantages, such as process simplicity and flexibility, and production of various organic chemicals depending on the type of catalyst employed.4-6 Since Y. Hori’s initial reports on CO2 electroreduction in the 1980s, the number of publications have dramatically increased, especially in the last decade.7 As mentioned above, electrocatalysis on different metal and metal oxide electrodes results to the production of various oxygenates. For example, Cu electrodes can transform CO2 into C1 and C2 hydrocarbons (e.g. methane (CH4) and ethylene (C2H4)).6-8 Carbon monoxide (CO), however, is mainly produced on Au and Ag.6,9,10 On the other hand, CO2 can also be used to produce liquid chemicals, such as the case of formate (HCOO-) electrosynthesis on Hg, Pb, and Sn.6,11-14 In fact, several review papers have carefully organized the available literature on this subject, in addition to describing the possible reaction pathways for different metals.1-6 For instance, Azuma et al. summarized the product distributions of metal catalysts using the periodic table.5 An updated classification has also included metal alloys, metal complexes, and metal oxides.2 In contrast to metal catalysts, however, there have been far fewer reports on the use of metal oxides in CO2 reduction.15-24 The role of metal oxides whether as catalysts for the


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formation of oxygenates or as precursor for the fabrication of well-structured catalysts, still remains unclear. Notably, Kanan’s group published several reports on the metal oxide effect in CO2 reduction.21-23 For example, the CO2 reduction efficiency highly depends on the presence of SnOx, with Sn/SnOx thin film electrodes catalyzing the formation of CO and HCOOH as main reaction products.22 Meanwhile, initial SnOx thickness in an electrode prepared by reducing a SnOx layer after 20 min of pre-electrolysis affects the product selectivity.24 Formate or formic acid as highly valuable oxygenates from CO2 have been receiving great attention because of their versatile use in various applications (e.g. direct formic acid fuel cells, leather, textile, and food industries) and low energy requirement during electrosynthesis in inexpensive electrodes.4,12-14 While SnO2 exhibits good catalytic activity for HCOO- production, literature reports cite difficulty in maintaining the SnO2 phase in the highly reducing conditions of CO2 electrolysis, a problem common among metal oxide catalysts.17-19,21-24 CO2 electrolysis in alkaline conditions could offer the possibility of maintaining the oxide stability.25 Hence, we explored the optimum E-pH conditions to selectively produce HCOO- at a high faradaic efficiency and electrode stability. In addition, the pH and potential effects on the reactant and product distributions were demonstrated as well.

2. Experimental Method The Sn/SnO2 nanocatalysts were deposited on the electrode by spraying. The catalyst ink consisted of SnO2 nanopowder (Sigma-Aldrich, >99 %, average particle size of 100 nm, 3 mg cm-2 loading) and Sn nanopowder (Sigma Aldrich, > 99 %, average particle size of 150 nm, 3 mg cm-2 loading) with 30 wt. % Nafion solution (Sigma-Aldrich, 10wt %) and 5 mL of 1-propanol (Aldrich, 99 %) to disperse the components. The mixture was mixed ultrasonically for 20 min


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and subsequently applied onto a gas diffusion layer with an area of 4 cm2 (GDL, 10 BC, SGL). The spraying process was performed using a spraying gun (Fuso Seiko Co., available nozzle diameter of 0.35 mm, storage volume of 8 cm3) while the GDL was placed on a hot plate at the drying temperature ~70 °C. All electrochemical experiments were carried out using a threeelectrode assembly, with a commercial Hg/HgO reference electrode and the potentials were converted to the RHE scale. Pt plate (1 cm2) was used as the counter electrode. The overall CO2 electrolysis set up is described in Figure S1. Electrochemical reduction of CO2 was performed in chronoamperometry (CA) mode using a potentiostat/galvanostat (PGSTAT-302N, Autolab). The electrolyte was made from deionized water and sodium hydroxide (NaOH, >97%, Daejung) with a catholyte volume of 90 mL. For CO2 electrolysis, high purity CO2 gas (99.999 %) was bubbled into the electrolyte at a flow rate of 40 mL min-1. Linear sweep voltammetry (LSV) measurements were obtained to investigate electrochemical behaviour of the electrodes in CO2-dissolved NaOH solutions at three pH levels ( 8.42, 10.2 and 11.72), as well as in a pure 0.5 M NaOH solution. The scan rate was fixed at 20 mV sec-1. In order to explore the optimal condition for CO2 electroreduction, CA was carried out at fixed potentials between -0.4 V and -1.6 V (vs. RHE), for 20 min at each potential. Subsequently, long-term stability test was evaluated at -0.6 V (vs. RHE) in the weakly alkaline pH of 10.2 over 5 h on both SnO2 and Sn electrodes. X-ray diffraction (XRD, Rigaku Miniflex-II) studies were carried out to observe the phase and crystallinity changes before and after CO2 electrolysis. To prepare the electrodes for XRD analysis, the used electrodes were rinsed with deionized (DI) water and dried instantly by hot-air blowing to minimize the formation of tin hydroxides and oxides on the electrode surface. We also verified the surface state by using X-ray photoelectron spectroscopy (XPS, VG Multilab


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2000). The liquid phase products generated from the electrolysis were analyzed using UVspectroscopy (UV-1800, Shimadzu) and high-performance liquid chromatography (HPLC, Alliance 2690, Waters) with Shodex RSpak KC-G and KC-811 columns (Figure S2). In UV analysis, HCOO- peak was observed at the wavelength of 230 nm, which is shifted from its wellknown reference position of 210 nm due to the Bathochromic shift by solvent effect. The HPLC analysis also confirmed that HCOO- was the only liquid product generated from the CO2 electrolysis (Figure S2). Perchloric acid (HClO4) was used as the mobile phase at a flow rate of 1 mL min-1 and a column temperature of 25 °C. Furthermore, gas-phase reaction products were measured by gas chromatography (GC, Agilent 7890A, Agilent Techonologies) equipped with thermal conductivity detector (TCD) (Figure S3). Carboxen 1006 PLOT column (Superico) was used to detect the hydrogen (H2) under the condition of nitrogen (N2) flowing as a carrier gas at 1.5 mL min-1.

3. Results and Discussion Highly sustained synthesis of HCOO- on SnO2 electrodes, at a high molar flow rate, is challenging due to the in-situ reduction of SnO2 into metallic Sn, resulting to a decrease in the faradaic efficiency. Solving this challenge opens up three main questions to which we propose to answer in this work. First, which carbon species serve as the reactant in the CO2 electroreduction on SnO2 to produce HCOO-? Second, at which pH to operate CO2 electroreduction that will maintain the SnO2 structure? Lastly, is SnO2 still superior over Sn in the selective production of HCOO- in alkaline media? Figure 1 shows the combined Pourbaix diagram indicating the overlap region between the dissolved species of CO2 and the phases of Sn-SnO2 in aqueous solutions at different pH and


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potentials.25 The hydroxides such as Sn(OH)2 and Sn(OH)4 are thermodynamically less stable than their corresponding oxides. In addition, the hydroxides have a tendency to change into their oxide form. Hence, as described in the Figure 1, we considered solely solid phase materials such as metallic Sn, and the anhydrous oxides SnO and SnO2.25 SnO is not depicted on the Pourbaix diagram because of its instability as compared to Sn and SnO2. In fact, metallic Sn is directly oxidized to SnO2 in weakly acidic and neutral media. The thermodynamic redox potentials of Sn/SnO and Sn/SnO2, which are close to one another, might explain why it is difficult to consider the presence or predominance of SnO in the aqueous system for the electrolysis. Inferring from the E-pH diagram, one can hypothesize that SnO2 could catalyze the electroreduction of CO2 to HCOO- while avoiding the reduction of the electrode into metallic Sn. A previous work using Sn electrode performed a pre-electrolysis for 20 min in 0.1 M KHCO3 to reduce the SnOx electrodes.24 In the experiment above, reduction of SnO2 to Sn occurred in the CO2-containing 0.1 M KHCO3 solution (pH=8.3), consistent with the combined E-pH diagram. It was proposed that the transformation of the metal oxides follows:

SnO + 2H+ + 2e- → Sn + H2O

(1)

SnO2 + 4H+ + 4e- → Sn + 2H2O

(2)

Three alkaline pH levels were selected as conditions to study the behaviour of SnO2 electrodes in CO2–bubbled 0.5 M NaOH solutions: pH at 8.42, 10.2, and 11.72. As known in the carbonate equilibria in aqueous systems, HCO3- species almost solely exist at pH=8.42, whereas CO32- species are predominant at pH=11.72.26 Meanwhile, at pH=10.2, both species have similar


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concentrations. It should be noted as well that these three pH levels are well within the region of interest, where SnO2 could be maintained during CO2 electrolysis. At each pH, linear sweep voltammetry (LSV) experiment was performed to study the CO2 electrolysis on SnO2 electrodes. As seen in Figure 2, the current densities increased with the decrease in pH. This is mainly due to the hydrogen evolution reaction (HER) and the possible reduction of SnO2 to Sn.27,28 In addition, a reduction peak was observed at around -0.6 V (vs. RHE) at the three pH levels. Consistent with previous studies, the peaks can be attributed to the electroreduction of carbonate species in the aqueous solution.27,28 Further proof is provided by the disappearance of a reduction peak without CO2 in the system, as shown in the inset of Figure 2. The reactive molecules during CO2 reduction in SnO2 electrodes, however, are the carbonate species – and not necessarily CO2. When CO2 dissolves in an aqueous solution, the following equilibria are established: CO2(g) ↔ CO2(aq)

(3)

CO2(aq) + H2O ↔ H2CO3

(4)

H2CO3 ↔ HCO3- + H+ pK1 = 6.35

(5)

HCO3- ↔ CO32- + H+

(6)

pK2 = 10.33

At the near neutral pH range between 6 and 9, which is generally employed for CO2 electroreduction, H2CO3/HCO3- are the predominant species. In alkaline media, however, the molar fraction of H2CO3 is almost zero. While the reactant that is mainly converted into larger carbonaceous molecules via electrochemical reduction is still under debate, most evidences point out to HCO3- as the reactant. For instance, Hori et al. proposed that HCOOH forms directly from CO2(aq), through the dissociation of HCO3-.29 This is supported by another study where higher faradaic efficiency for HCOO- was demonstrated at pH=8, arising from the relatively fast


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dissociation of HCO3- to CO2.30 In addition, Innocent and co-workers reported that CO2 reduction in Pb electrodes effectively occurs at the pH where HCO3- predominantly exists and CO32- was not reactive at the electrode surface.27 Recent results from Koper’s group show that the direct HCO3- reduction to HCOO-, which is claimed to occur in palladium (Pd), also happens on Cu electrodes.28,31,32 UV-spectroscopy was performed to detect possible liquid phase products that are formed during CO2 electrolysis. A constant potential was applied at -0.6 V (vs. RHE), the potential where the reduction peaks appeared in the LSVs. Only formate was confirmed as the liquid phase reaction product, as displayed in the Figure S2. Moreover, the relative product concentrations correspond well with the magnitude of the currents in the LSVs in Figure 2. At the more alkaline conditions where CO2 is almost non-existent, HCOO- formation indicates the direct reduction of HCO3-. An additional proof is provided by the production of HCOO- in a pure 0.5 M NaHCO3 solution without CO2-purging (Figure S4). As the formate was generated at the same reduction potential and pH in the experiment earlier, direct transformation of HCOO- from HCO3- is further verified.28,31,32 On the other hand, CO2 electrolysis also produced CO at pH=8.42. At higher pH, however, gas chromatography (GC) analysis did not observe any gaseous organic products, contrary to previous results showing CO gas produced on Sn/SnOx electrodes.22,24 Recent studies reported the pH dependence of product distribution in CO2 reduction.33,34 For example, scanning electrochemical miscroscopy (SECM) investigations of CO2 reduction on Au electrodes show that HCOO- comes solely from the direct reduction of HCO3- and CO forms only via CO2 reduction.33 Another report explored the dependence of the molar ratio between CO and HCOOgenerated on Sn electrodes as a function of pH, with the ratio ranging from 1 to 0.15 over the pH


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range of 2.9 to 7.8.34 These studies demonstrate that electrolysis in alkaline condition could selectively enhance the synthesis of HCOO- over CO on SnO2 electrodes. A catalyst’s stability in an actual reaction condition is as important as its activity and selectivity. Hence, the phase changes and crystallinity of the SnO2 electrodes were studied via ex-situ X-ray diffraction (XRD) before and after CO2 electrolysis. As shown in Figure 3a, SnO2 was partially reduced into metallic Sn at pH=8.42. In contrast, at higher pH values of pH=10.2 and pH=11.72, SnO2 remained intact and Sn diffraction peaks were not observed. While the pH difference seems small, it could significantly affect CO2 electrolysis due to the drastic change in the equilibrium concentrations of the carbonate species in the pH region considered.26 Figure 3b shows the Sn 3d5/2 XPS spectra indicating the oxidation state of the electrode surface, with the peaks appearing at 486.8 eV and 485.0 eV corresponding to SnO2 and metallic Sn.24,35 The peaks also demonstrated that the Sn oxide layer on the electrode surface was not reduced to Sn during the electrolysis in alkaline media. As seen in Figure 3b-(v), the SnO2 peak was enriched after the commercial Sn powder was exposed in air. Nevertheless, the metallic Sn peak is still visible in the deconvoluted spectrum. The in-situ reduction of SnO2 to Sn at pH=8.42 was also observed, as shown by the presence of a low binding energy peak. This is consistent with the results in XRD where the reduction of SnO2 to Sn after CO2 electrolysis was also observed. To facilitate easier comparison between the SnO2 electrocatalytic behaviour at the three different pH values considered, Table 1 lists the related observations during CO2 electrolysis. At pH=8.42, where HCO3- predominantly exists, the highest current density and HCOO- production rate were observed. Ex-situ X-ray diffractograms, however, confirmed the instability of SnO2 at this pH. Furthermore, a large fraction of the current was used for the HER, leading to a large


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efficiency loss when large oxygenates are the target products. On the contrary, SnO2 phase was maintained at the two higher pH values even under the applied reducing potential, which is an indication of the stability of the electrode. Hence, subsequent studies only focused on the two more alkaline pH levels. As displayed in Table 1, the relative concentration of CO32- represents as the foremost identifying factor between the two alkaline pH levels. It must be noted, however, that CO32- is considered as non-reactive species in a CO2 reduction system. Nevertheless, it would interesting to know if CO32- could act as an effective reservoir, where the electroactive carbonate species can be derived following the reverse reaction in Equation 6. When considering a system where CO2 is in equilibrium with a 0.5 M NaOH solution at pH=11.72, the total concentration of dissolved carbon-containing species is estimated to be 2.48 × 10-1 M, while the molar fraction of CO32- is 0.94 (See detailed computations in SI). Assuming that the remainder is HCO3- only, the total amount of reactive species available for conversion to various products during CO2 reduction is only around 1.48 × 10-2 M. Therefore, as HCO3- is continuously consumed during CO2 reduction, additional HCO3- might be produced by the CO32- reservoir. To test this theory, CO2 electrolysis was performed following the conditions above at pH=11.72, with the product samples collected every 20 min to measure HCOO- production before and after CO2-bubbling. As seen in Figure 4, production of HCOO- significantly dropped with time, signifying mass transfer limitation at the electrode surface. While the carbonate equilibria renormalizes to form HCOO-, depletion of HCOO- at the electrode surface during CO2 reduction occurs at a notably faster rate. After CO2 was bubbled at the 60 min mark, however, the faradaic efficency for HCOO- and the total cathodic current density rose abruptly. This should provide ample evidence that at the high pH where CO32- predominates, CO2 electrolysis


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at a high faradaic efficiency can only effectively occur in the presence of CO2 in the solution, albeit at a lower production rate than at a lower pH. Hence, we suggest that the optimal pH is at or near pH=10.2, where the condition provides enough alkalinity to prevent reduction of SnO2 and allows higher production rates of HCOO-. Subsequently, the optimal potential for efficient CO2 electrolysis at pH=10.2 was explored from -0.4 V to -1.6 V (vs. RHE). As shown in Figure S5, the highest faradaic efficiency for HCOO- formation was observed at -0.6 V. Applying a more negative potential starting from 0.4 V increased the faradaic efficiency up to -0.6 V only. Beyond -0.6 V, higher overpotential no longer led to increased faradaic efficiency. This is in agreement with the observed cathodic peak at around -0.6 V in the LSVs. The location of the highest faradaic efficiency for HCOOformation at the reduction peak potential has been reported as well in previous studies.27,34 This is presumably due to the acceleration of the competitive HER at more negative potentials. In addition, the CO2 reduction was also performed using metallic Sn catalysts, in order to provide comparison with SnO2 catalysts. SnO2 exhibited better catalytic performance over Sn, as shown by higher faradaic efficiencies at all applied potentials. The difference in activity between SnO2 and Sn nanocatalysts most likely arises from their inherent catalytic activities and not from the physical properties. Both catalysts have also similar morphologies before and after CO2 electrolysis (Figure S6). Although the Brunauer-Emmett-Teller (BET) surface area of SnO2 is twice that of Sn, its superior catalytic activity is not due to this difference. To demonstrate this, SnO2 was fabricated by annealing the Sn nanocatayst at 150 °C for 10 h. As shown in Table S1, the annealed SnO2–based electrode showed enhanced catalytic activity for CO2 reduction even with the decrease in surface area.


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The logarithmic partial current density ratios of HCOO- and H2 on SnO2 and Sn electrodes, a measure of product selectivity, were plotted with the respect to the applied potential (see Figure 5a). Consistent with the observations earlier, SnO2 generally showed better selectivity for the production of HCOO- over Sn. Both electrodes demonstrated high production rates at -0.6 V, with decreasing current densities from -0.6 V to -1.0 V. As the potential becomes more negative than -1.0 V, however, there is a strong likelihood that SnO2 is reduced to Sn. This is evident from the slightly increasing selectivity, despite the decreasing faradaic efficiency for HCOO- formation at this potential region. It must be noted that the partial current density for electrode phase change was excluded when calculating the product ratios. While the reduction of SnO2 to Sn at more negative potentials was not detected in ex-situ XRD analysis, the results above should provide a strong warning about operating the CO2 electrolysis at high overpotentials even in alkaline media. In order to demonstrate the feasibility of sustained production of HCOO- from CO2 on SnO2 electrodes in the optimal E-pH condition, long term CO2 electrolysis was performed at 0.6 V and pH=10.2. The absorbance points of samples obtained during the electrolysis is listed in Table S2. The liquid products were collected and analysed every 60 min and HPLC was also employed to verify the concentration of HCOO- produced in the electrolyte solution. As seen in Figure 5b, a high faradaic efficiency of ~53% was maintained even after 5 hr on SnO2 electrode. As shown in Figure S6, we did not observe significant morphological change such as agglomeration that could lead to the degradation of catalytic activity. In addition, the material loss was also insignificant in our electrolytic system. Although our system showed superior catalytic stability over 5 h, it is believed that the slight decrease of activity can be attributed to the formation of intermetallic compounds.36,37 Quite recently, Anawati et al. reported that the


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formation of intermetallic compound, which occurs during CO2 reduction, leads to both cathodic deactivation and material loss. As already established in earlier results, metallic Sn electrode showed relatively lower catalytic performance. In view of the oxide reduction commonly encountered in CO2 electroreduction catalysts, electrolytic operation at the optimal E-pH condition avoids the electrochemical reduction of SnO2 to Sn. This is supported by the absence of metallic Sn peaks in the ex-situ X-ray diffractograms of the catalyst after the 5 hr of operation (Figure S7). In summary, we have demonstrated that SnO2 is the better CO2 reduction catalyst for the production of HCOO- in comparison with Sn, but the CO2 electrolysis must be performed at an optimal E-pH in order: (1) to selectively produce HCOO- over CO formation (2) to allow high production rates of HCOO- arising from high faradaic efficiency and reduction currents, and (3) to intrinsically provide stability to the SnO2 catalyst for long-term operation.

4. Conclusion We have investigated the effects of pH and applied potential on the catalytic performance of SnO2 electrodes in terms of activity, selectivity, and stability. Answering the questions posed earlier, we demonstrate that HCO3- species serve as the reactant reservoir for HCOO- formation. Meanwhile, operation in highly alkaline pH with predominantly CO32- species is severely limited by the short supply of reactants at the electrode surface. The superior electrocatalytic behaviour of SnO2 in comparison to Sn was also confirmed, in agreement with earlier studies done in neutral pH. Additionally, SnO2 showed remarkable stability at the optimal E-pH region due to the prevention of in-situ electrode reduction. These items highlight the possibility of operating CO2 electrolysis that are using metal oxide electrodes in weak alkaline conditions in order to enhance the oxide structure stability, a known issue among metal oxide catalysts. While most


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studies have so far focused on metal catalysts in neutral or near-neutral pH, the use of metal oxide at higher pH could allow for sustained higher faradaic efficiencies. This approach is widely applicable to other metal-metal oxide systems, especially nanostructured CO2 reduction catalysts, and could hopefully lead to practical production of valuable and carbon-neutral chemicals from CO2 via renewable energy.

Supporting Information: Schematic diagram of the CO2 electrolysis set-up, UV spectra of the electrolysis products at different pH levels, HPLC results for the liquid phase product of CO2 electrolysis, GC results for the gas phase product of CO2 electrolysis, UV spectra of electrolysis products with and without CO-bubbling, HCOO- production at different applied potentials in SnO2 and Sn electrodes, X-ray diffractograms of the SnO2 nanocatalysts before and after CO2 electrolysis at the optimal E-pH condition. AUTHOR INFORMATION Corresponding Author: *Tel: +82-62-715-2440. Fax: +82-62-715-2434. E-mail: [email protected] Notes: The authors declare no competing financial interest.

ACKNOWLEDGMENTS This research was supported by Basic Science Research Program through the National Research Foundation

of

Korea

(NRF)

funded

by

the

Ministry

of

Education

(NRF-

2013R1A1A2A10063010).


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Figure 1. Combined potential-pH (E-pH) equilibrium diagram of the tin-water system considering exclusively the anhydrous oxide phases and the carbonate-water system with the relative dominance of each carbonate species. The encircled area indicates the overlap region where a stable SnO2 phase could catalyze HCOO- electrosynthesis from CO2. The E-pH diagram was redrawn from individual E-pH diagrams in Ref. [25].


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Figure 2. Comparison of linear sweep voltammetry (LSV) curves on the SnO2 electrodes in CO2-dissolved 0.5 M NaOH solutions at (a) pH=11.72, (b) pH=10.2 and (c) pH=8.42. Inset voltammogram was obtained in a pure 0.5 M NaOH with Ar gas-purging at pH=13.7. The scan rate was fixed at 20 mV sec-1.


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Figure 3. (a) XRD patterns and (b) XPS spectra of SnO2 electrodes: (i) as-prepared SnO2 and SnO2 electrodes after CO2 electrolysis at (ii) pH=11.72, (iii) pH=10.2, (iv) pH=8.42, and (v) commercial Sn powder as reference.


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Figure 4. Comparison of total current density and faradaic efficiency for HCOO- on SnO2 electrode at initial pH 11.72 before (current density : black-dotted line, faradaic efficiency : ■) and after (current density: red-solid line, faradaic efficiency : ●) the addition of CO2 gas


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Figure 5. (a) Product selectivity as represented by the ratio of faradaic current densities for HCOO- and H2 production in log function form on (i) SnO2 and (ii) Sn electrodes. (b) Total current density profiles during the electrolytic process on (i) SnO2 (red-solid line) and (ii) Sn (black-dotted line). Faradaic efficiencies for HCOO- on SnO2 (●) and Sn (■) were measured every 60 min.


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Table 1. Comparison of the relevant parameters during alkaline CO2 reduction on SnO2 electrodes at the pH values considered.


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Alkaline CO2 Electrolysis towards Selective and Continuous HCOOProduction over SnO2 Nanocatalysts Seunghwa Lee,† Joey D. Ocon,† Young-il Son,†§ Jaeyoung Lee*,†,‡ †

Electrochemical Reaction and Technology Laboratory (ERTL), School of Environmental Science and Technology, ‡Ertl Center for Electrochemistry and Catalysis, RISE, Gwangju Institute of Science and Technology (GIST), Gwangju 500-712, South Korea § Korea Environmental Industry & Technology, Seoul 122-706, South Korea.

Table of Contents – Graphic

While the electrochemical conversion of CO2 into high value chemicals in metal oxides offers high activity, neutral pH CO2 electrolysis results to in-situ reduction of SnO2 to metallic Sn. We demonstrate that shifting the reaction condition towards an optimal potential-pH region leads to a more selective formate production and stable SnO2 phase.


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Alkaline CO2 Electrolysis toward Selective and Continuous HCOO

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