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The impact of nitrogen oxides on electrochemical carbon dioxide reduction


The electroreduction of carbon dioxide offers a promising avenue to produce valuable fuels and chemicals using greenhouse gas carbon dioxide as the carbon feedstock. Because industrial carbon dioxide point sources often contain numerous contaminants, such as nitrogen oxides, understanding the potential impact of contaminants on carbon dioxide electrolysis is crucial for practical applications. Herein, we investigate the impact of various nitrogen oxides, including nitric oxide, nitrogen dioxide, and nitrous oxide, on carbon dioxide electroreduction on three model electrocatalysts (i.e., copper, silver, and tin). We demonstrate that the presence of nitrogen oxides (up to 0.83%) in the carbon dioxide feed leads to a considerable Faradaic efficiency loss in carbon dioxide electroreduction, which is caused by the preferential electroreduction of nitrogen oxides over carbon dioxide. The primary products of nitrogen oxides electroreduction include nitrous oxide, nitrogen, hydroxylamine, and ammonia. Despite the loss in Faradaic efficiency, the electrocatalysts exhibit similar carbon dioxide reduction performances once a pure carbon dioxide feed is restored, indicating a negligible long-term impact of nitrogen oxides on the catalytic properties of the model catalysts.


The electrochemical CO2 reduction (CO2RR) provides a promising, sustainable avenue to generate value-added fuels and chemicals from greenhouse gas CO21,2. Depending on the choice of electrocatalyst, CO2 can be converted into a variety of single-carbon (C1; e.g., carbon monoxide, formic acid, methanol, and methane) and multi-carbon (C2+; e.g., ethylene, ethanol, acetate, and n-propanol) products with tremendous market potentials3,4,5,6,7,8. While CO2RR is being actively studied, most studies are conducted using highly pure CO2 feed9,10. For commercial applications, the most commonly available CO2 sources are industrial point sources, such as chemical and power plants;11 however, CO2 gas emitted from these sources often contain a variety of contaminants, such as sulfur oxides (SOx), nitrogen oxides (NOx), O2, and volatile organic compounds (VOC) (Fig. 1a)12,13,14. Therefore, there is an urgent need to understand the potential impact of common contaminants in industrial CO2 sources on the catalyst properties in CO2RR.

Fig. 1: CO2 electrolysis technology using industrial CO2 point sources.
figure 1

(a) Schematics of CO2 electrolysis with CO2 stream obtained from point sources containing impurities such as nitrogen oxides (NOx), sulfur oxides (SOx), O2, and volatile organic compounds (VOC) and potential influence of impurities in CO2 electroreduction (CO2RR). (b) Standard potential vs. reversible hydrogen electrode (RHE) for CO2RR, hydrogen evolution reaction (HER), NO2 reduction (NO2RR), NO reduction (NORR), and N2O reduction (N2ORR). Detailed reactions are provided in Supplementary Table 1.

Gas impurity in CO2 can affect the performance of CO2RR electrocatalysts as we demonstrated in the case of SO215, where a trace amount of SO2 in the feed is sufficient to alter the product selectivity of Cu catalyst substantially. The potential impacts of impurity include lowering Faradaic efficiency (FE; i.e., number of electrons transferred to desired products divided by the total number of electrons passed in the system) due to competing reactions of impurity over CO2, altering the property of the catalyst by incorporating into the catalyst and/or support, and adsorbing on the catalyst surface to physically block the active sites (Fig. 1a). To date, there are only a few studies focusing on understanding how the presence of contaminants influences the behavior of electrocatalysts under CO2RR conditions15,16,17,18,19,20,21. For example, NOx is one of the major contaminants present in industrial CO2 point sources with a typical concentration of 1000 ppm12,13,14. The NOx contaminants typically consist of 90–95% nitric oxide (NO) and 5–10% nitrogen dioxide (NO2)22. Additionally, nitrous oxide (N2O) is also a common byproduct formed in the NOx removal process23, which has a relatively low reactivity in comparison to other NOx. A previous study has shown that 200 ppm of NO has a negligible influence on Cu catalysts in CO2RR in a conventional batch cell18. Furthermore, less than or equal to 1667 ppm of NO2 has shown to be either beneficial or neutral, and greater than 1667 ppm of NO2 has shown to be detrimental in CO2RR, mainly due to a reduction in pH of the electrolyte, also on Cu catalysts in a conventional batch cell19. However, the behavior of various NOx impurities in CO2RR at industrially relevant high current densities (>100 mA cm−2) has not been explored yet.

In this work, we investigate the influence of NOx (i.e., NO, NO2, and N2O) in CO2RR using a three-compartment flow cell. Three model electrocatalysts, including copper (Cu), silver (Ag), and tin (Sn), are selected to represent the most studied catalysts for C2+ products, carbon monoxide (CO), and formate, respectively. Most NOx contaminants in the CO2 feed significantly reduce the CO2RR FE because the electrochemical reduction of NOx occurs at much more positive potentials than CO2RR (Fig. 1b). NO and NO2 impurities have more severe impacts on CO2RR FE than N2O, likely due to the greater number of electrons required in the NOx reactions. Despite the loss of CO2RR FE, none of the three catalysts exhibits a significant change of product selectivity after removing the NOx impurity from the CO2 feed. Moreover, we employ gas chromatography (GC), spectrophotometry, and flow electrochemical mass spectrometry (FEMS) to analyze the products of electroreduction of NO, the dominant component of NOx in industrial point sources, in which the major products are ammonia (NH3), hydroxylamine (NH2OH), N2, and N2O. Investigation of the effect of different concentrations of NO in CO2RR shows that NOx at typical concentrations in flue gases is compatible with CO2RR.

Results and Discussion

Electrodes were prepared by loading commercial Cu, Ag, and Sn particles on a gas diffusion layer (GDL), a microporous carbon paper which provides mechanical support, electrical conductivity, and hydrophobicity. Scanning electron microscopy (SEM) images of the as-prepared electrodes confirm a uniform deposition of metal nanoparticles on GDL, covering the majority of the GDL surface (Supplementary Fig. 1). Electrochemical experiment was performed in a three-compartment flow cell, in which CO2 gas is directly fed to the electrode-electrolyte interface, enabling CO2RR at high current densities (Supplementary Fig. 2). NOx impurities were mixed with CO2 gas feed prior to entering the flow cell. As the concentration of NOx in typical exhaust streams may be as high as ~3,000 ppm (i.e., 0.3 vol. %)13, conservative streams of 83.3% CO2, 15.87% Ar, and 0.83% NOx were used for most studies. To keep the CO2 partial pressure constant during the introduction of NOx, which contains Ar, CO2 partial pressure was maintained at 0.833 bar throughout the study by using a mixture of 83.3% CO2 and 16.7% Ar when NOx was not introduced.

Impact of NOx impurities on CO2 electroreduction

The influence of NO in CO2RR on Cu, Ag, and Sn catalysts was first evaluated at a constant current density of 100 mA cm−2 (Fig. 2a–c). The CO2RR experiment was performed by switching the gas feed from 83.3% CO2 and 16.7% Ar (0–0.5 h) to 83.3% CO2, 15.87% Ar, and 0.83% NO (green region; 0.5–1 h) and back to 83.3% CO2 and 16.7% Ar (1–3 h). With 83.3% CO2 and 16.7% Ar, before exposure to NO, Cu catalyst produced a wide range of C1 (i.e., methane, CO, and formate) and C2+ (i.e., ethylene, ethanol, acetate, and propanol) products. In the cases of Ag and Sn catalysts, the major products were CO and formate, respectively. The observed CO2RR selectivity of the Cu, Ag, and Sn catalysts was consistent with the previous reports6,7.

Fig. 2: CO2 electroreduction performance in the presence of NO.
figure 2

Faradaic efficiency and applied potential vs. time on (a) Cu, (b) Ag, and (c) Sn catalysts at a constant current density of 100 mA cm−2 in 1 M KHCO3 for 3 h. Gas feeds were 83.3% CO2 and 16.7% Ar, and 83.3% CO2, 15.87% Ar, and 0.83% NO (green). 0.83% NO was introduced at 0.5 h for 0.5 h. Corresponding Faradaic efficiencies are provided in Supplementary Tables 24. (d) Effect of different concentrations of NO in CO2 electroreduction on Cu, Ag, and Sn catalysts. 0.083% and 0.0083% represent the typical NOx concentrations in flue gases and flue gases after NOx removal processes, respectively. Corresponding Faradaic efficiencies are provided in Supplementary Table 6. Error bars represent the standard deviation of three independent measurements.

When 0.83% NO was introduced at t = 0.5 h, the total CO2RR FE decreased noticeably on all three catalysts (Fig. 2a–c). On average, the losses in CO2RR FE accounted for 33.9, 29.6, and 27.9% on Cu, Ag, and Sn, respectively (Fig. 2d), which is likely due to the preferential reduction of NO over CO2. Assuming NO is fully converted to NH3, conversions of NO during CO2RR are between 48% and 60% (Supplementary Table 5). As shown in Fig. 1b, the standard potentials of NORR are much more positive than those of CO2RR. For instance, the standard potential of NORR to N2 is 1.68 V vs. RHE, while the standard potentials of CO2RR are between −0.250 and 0.169 V vs. RHE. Cyclic voltammetry (CV) measurements under CO2 with 0.83% NO also confirmed that NORR is more favorable than CO2RR on Cu, Ag, and Sn catalysts (Supplementary Fig. 3). On all three catalysts, onset potentials and cathodic currents shifted to more positive potentials when 0.83% NO was introduced to the CO2 stream. CV measurements under different concentrations of NO in Ar also confirmed more positive onset potentials of NORR than CO2RR and showed that NORR at 0.83% NO is mass transport limited (Supplementary Fig. 4). After restoring 83.3% CO2 and 16.7% Ar, the CO2RR performance and the total CO2RR FE on all three catalysts quickly recovered and were stable for additional 2 h of electrolysis. No obvious change in selectivity was observed for any of the three catalysts, suggesting that the exposure to NO did not alter the catalyst property in any significant way. There is a slight increase in H2 FE over time (Fig. 2a–c), but it is likely due to the slow flooding of the electrode (Supplementary Fig. 5)24.

To obtain insight on the influence of NOx in CO2RR at typical concentrations of NOx in point sources, we evaluated the effect of 0.083% and 0.0083% NO, representing the typical NOx concentrations in flue gases and flue gases after NOx removal processes22, respectively, in CO2RR (Fig. 2d). Although the losses in FE at 0.83% NO were detrimental, the effect of NO was less severe at 0.083%, with less than 5% losses in FE, and negligible at 0.0083% NO. Therefore, NO at typical concentrations of NOx in flue gases is compatible with CO2RR, although complete removal of NOx is desired to maximize CO2RR FE.

NO2 is another major contaminant in industrial CO2 point sources (5–10% of NOx), and a substantial amount of N2O may also be formed as a byproduct during the NOx removal process22,23. Thus, we further investigated the influence of NO2 and N2O in CO2RR on Cu, Ag, and Sn catalysts following the similar experimental procedure to the NO experiment. When 0.83% NO2 was introduced at t = 0.5 h (yellow region), the CO2RR FE decreased on all three catalysts (Fig. 3a). The decrease in the total FEs were 30.8, 25.6, and 22.9% on Cu, Ag, and Sn catalysts, respectively. Similarly, when 0.83% N2O was introduced (blue region), the total CO2RR FE decreased by 11.4, 10.2, and 1.4% on Cu, Ag, and Sn catalysts, respectively (Fig. 3b). Distinct from Cu and Ag catalysts, Sn catalyst did not show a significant loss of the CO2RR FE in the presence of N2O, which is likely due to the poor activity of Sn for N2ORR25. Sn catalyst maintained a high CO2RR FE over the course of 3 h of electrolysis, suggesting the resistive feature of Sn catalyst to N2O impurity. As shown in Fig. 1b, standard potentials of NO2 and N2O are also more positive than those of CO2RR, and therefore, we attribute the loss of the CO2RR FE to the preferential reduction of NO2 and N2O over CO2, which is further supported by the CV study (Supplementary Figs. 811). When a pure CO2 feed was restored, the total CO2RR FE on all three catalysts quickly recovered, suggesting that the exposure of NO2 and N2O does not affect the property of the catalysts.

Fig. 3: CO2 electroreduction performance in the presence of NO2 and N2O.
figure 3

CO2 electroreduction Faradaic efficiency, excluding hydrogen Faradaic efficiency, vs. time with the introduction of (a) 0.83% NO2 (yellow) and (b) 0.83% N2O (blue) on Cu, Ag, and Sn catalysts at a constant current density of 100 mA cm−2 in 1 M KHCO3 for 3 h. Gas feeds were 83.3% CO2 and 16.7% Ar, and 83.3% CO2 and 15.87% Ar with 0.83% NO2 or 0.83% N2O. NO2 and N2O were introduced at t = 0.5 h for 0.5 h. Corresponding Faradaic efficiencies are provided in Supplementary Figs. 6 and 7, and Supplementary Tables 712. (c) Loss in Faradaic efficiency during CO2 electroreduction from the introduction of 0.83% NO, 0.83% NO2, and 0.83% N2O on Cu, Ag, and Sn catalysts. Corresponding Faradaic efficiencies are provided in Supplementary Table 13. Error bars represent the standard deviation of three independent measurements.

A comparison of the losses in FE due to the various NOx impurities is presented in Fig. 3c. NO and NO2 show greater losses in FE than N2O on all three catalysts, likely due to the greater number of electrons required in the reactions. As will be discussed in the following section, the main products of NORR are NH3 and NH2OH, which require 5 and 3 electrons, respectively, while the main product of N2ORR is N2, which only requires 2 electrons. Given that all NOx readily reacts at the catalyst surface, the same amount of NO and NO2 consume more electrons than N2O, causing greater losses in CO2RR FE. Among all the catalysts, the Cu catalyst suffers the largest FE loss on all NOx impurities, followed by Ag and Sn catalysts. Indeed, Cu has been demonstrated as one of the more active metals for the electroreduction of NO26 and N2O25, in which Cu achieved high FE in N2ORR to N2 at relatively low overpotentials. The results suggest that Cu is an effective electrocatalyst for NOx reduction, which may be further explored in future studies.

Furthermore, pH was measured at the outlet of the electrolyzer at different time points (i.e., before, during, and after NOx introduction) to investigate the effect of NOx on the electrolyte pH (Supplementary Fig. 12). The measured pH shows that the presence of NO and N2O has a negligible effect on the pH, while the presence of NO2 slightly decreases the pH by 0.03. Although NO2 hydrolyzes to produce nitric acid and nitrous acid27, the effect in pH is very small, possibly due to the small amount of NO2 in the gas feed, rapid reaction of NO2 at the catalyst surface which prevents NO2 from penetrating to the bulk electrolyte, and a flowing electrolyte which is constantly replenished.

Identification of NOx reduction products

The electrochemical reduction products of NO, the major component of NOx in industrial point sources, were further investigated. As NH3, NH2OH, N2, and N2O have been suggested as the main products in NORR26,28,29, NH3 and NH2OH were detected via spectrophotometry (Supplementary Figs. 13 and 14), and N2 was detected via GC (Supplementary Fig. 15). We note that the concentration of N2O in the gas product stream was below the detection limit of GC, suggesting that N2O FE was below 2% FE on all three catalysts. As shown in Fig. 4a, NORR product selectivity varied among different catalysts. Cu primarily produced NH3 and N2, with no NH2OH, Ag produced a mixture of NORR products, and Sn primarily produced NH2OH. These observations are consistent with previous reports, in which Cu has been demonstrated as an effective catalyst for NORR to NH326, and Sn has been used as a dopant in Pt to shift the selectivity from NH3 to NH2OH in nitrate reduction30.

Fig. 4: Investigation of the NO electroreduction products.
figure 4

(a) Faradaic efficiency of NO electroreduction products produced during electrolysis with 83.3% CO2, 15.87% Ar, and 0.83% NO on Cu, Ag, and Sn catalysts at a constant current density of 100 mA cm−2 in 1 M KHCO3 for 3 h. Corresponding Faradaic efficiencies are provided in Supplementary Table 14. Error bars represent the standard deviation of three independent measurements. (b) Schematic of flow electrochemical mass spectrometry (FEMS) setup. (c) Measured current density vs. time, and deconvoluted MS signal vs. time for m/z = 2, m/z = 17, (d) m/z = 28, m/z = 30, and m/z = 44 from FEMS on Cu catalyst in 1 M KHCO3 with 0.83% NO in Ar. −0.90 V vs. RHE was applied for approximately 2 min starting at t = 1.5 min. NORR products have been deconvoluted using the mass spectra of individual products shown in Supplementary Fig. 17. Additional information is provided in the Methods section and Supplementary Figs. 18 and 19.

To further probe the formation of NORR products with greater sensitivity and determine the formation of N2O, we employed the FEMS (Fig. 4b and Supplementary Fig. 16), which allows us to continuously measure gas and volatile liquid products operando with a low detection limit and a short response time by continuously pulling products to the mass spectrometry (MS) near the surface of the electrodes (See Methods for more details). The MS probe was placed near the working electrode from the gas channel side, and the MS signals linked to possible products were tracked over time. We conducted the FEMS measurement on the Cu catalyst using 0.83% NO in Ar (in the absence of CO2), because the ionization of N2 (m/z = 28, 14) and N2O (m/z = 44, 30, 28, 14) produces the same fragments with CO2 and various CO2 reduction products31,32,33, complicating the reliable analysis of the NORR products (see Supplementary Note for more details). MS signals of the FEMS measurement under a continuous feed of 0.83% NO are presented in Fig. 4c, d. When a constant potential of −0.90 V vs. RHE was applied at t = 1.5 for ~2 min, MS signals of NO (m/z = 30) decreased while those of H2 (m/z = 2), NH3 (m/z = 17), N2 (m/z = 28) and N2O (m/z = 44) increased (Fig. 4c, d), indicating the consumption of NO and the formation of H2, NH3, N2, and N2O. The formation of NH3 and N2 detected by FEMS is in agreement with the results obtained from spectrophotometry and GC analysis, respectively. The production of N2O, which was difficult to measure via GC, was clearly observed in FEMS, suggesting that N2O is one of the NORR products. NH2OH was not detected in FEMS, because it is nonvolatile30. Similarly, FEMS results also suggest the formation of N2 and N2O on Ag and Sn catalysts (Supplementary Figs. 2025). However, the formation of NH3 was observed only on Ag and not on Sn, likely due to the small amount of NH3 produced on Sn. Collectively, NH3, NH2OH, N2, and N2O have been determined as the NORR products. The analysis of the NORR products further confirms that the loss in CO2RR FE is due to the preferential reduction of NO over CO2.

In the case of N2ORR, a substantial amount of N2 was quantified with a GC (Supplementary Fig. 26). While the losses of CO2RR FE were 11.4%, 10.2%, and 1.4% on Cu, Ag, and Sn catalysts, respectively, the amounts of N2 detected were 8.2%, 7.3%, and 0.5% of the total FE, respectively, accounting for the majority of the loss in the CO2RR FE. Small amount of N2 detected on Sn catalyst demonstrates the resistive nature of Sn catalyst in N2ORR.

Characterization of catalyst structures in the presence of NOx

X-ray photoelectron spectroscopy (XPS) measurements were conducted to reveal the influence of NOx on the surface electronic structure and the chemical environment of the catalysts. The samples were obtained at various points of the CO2RR experiment, including before exposure to NOx, after exposure to NOx, and at the end of 3-h electrolysis. As shown in Fig. 5a, the Cu and Sn electrodes before the exposure to NOx did not show any noticeable peak in N 1 s XPS measurements. In contrast, Ag showed two distinct peaks at 400.5 eV and 398.5 eV, which can be attributed to polyvinylpyrrolidone (PVP)34,35, a surfactant used in the nanoparticle synthesis. The XPS measurements obtained after the NO exposure (t = 1 h) exhibited new N 1 s peaks on Cu and Sn electrodes (Fig. 5b). The peaks at 401.4 eV, 400.2 eV, and 398.2 eV can be assigned to graphitic, pyrrolic, and pyridinic N, respectively36,37, suggesting that incorporated N atoms mainly interact with carbon in GDL rather than metal catalysts (metal nitride peaks typically observed near 397 eV)38,39. The XPS measurements obtained after 3-h electrolysis show that the N incorporated in the electrode surface was still intact after additional 2 h of CO2RR (Supplementary Fig. 27 and Supplementary Table 15), with the total amount of N in the Cu and Sn electrodes remaining relatively unchanged. In the cases of 0.83% NO2 and 0.83% N2O, the XPS measurements show similar N incorporation in GDL (Supplementary Figs. 28 and 29, and Supplementary Tables 16 and 17). Regarding the Ag electrode, the XPS investigation of N incorporation associated with NOx was largely limited by the presence of the PVP surfactant.

Fig. 5: Evaluation of the influence of NO on the catalyst structure.
figure 5

XPS measurements of Cu, Ag, and Sn electrodes (a) before (t = 0 h) and (b) after exposure to 0.83% NO (t = 1 h) during CO2 electrolysis. Corresponding XPS data is provided in Supplementary Fig. 27 and Supplementary Table 15. Cu K-edge (c) XANES and (d) EXAFS spectra of spent Cu catalyst after exposure to 0.83% NO during CO2 electrolysis. Cu foil, Cu2O, and CuO were used as references.

To further confirm the incorporation of N into GDL rather than the formation of metal nitrides, we increased the catalyst loading to 2.0 mg cm−2, which created a thick layer of catalyst on the GDL with much less exposure of GDL in the XPS measurement. After exposure to 0.83% NO during CO2 electrolysis, the N 1 s signal was not detected on the Cu and Sn electrodes with the increased catalyst loading, whereas the XPS measurements for the Ag electrode clearly shows the N 1 s signal, which is due to the presence of PVP on the surface of Ag catalyst (Supplementary Fig. 30). Conversely, when the same experiment was repeated with GDL without any catalyst, N species was still detected, confirming the incorporation of N into GDL (Supplementary Fig. 30). Experiments using NO2 and N2O show similar incorporation of N into GDL (Supplementary Fig. 30).

To probe the influence of NOx impurity on the oxidation state of the Cu catalyst, we conducted X-ray absorption spectroscopy (XAS) measurements using a customized XAS batch cell (Supplementary Fig. 31). Because of the toxicity of the NOx gases, we did not use the NOx gases directly at the synchrotron X-ray beamline but conducted XAS experiments with the electrodes taken out of the electrolyzer at 1 h (after exposure to NOx for 0.5 h) during CO2 + NOx experiments (Fig. 2a-c and Supplementary Figs. 6 and 7). The Cu K-edge X-ray absorption near-edge spectroscopy (XANES) spectra of the Cu catalyst after the NO exposure show a similar spectrum of the Cu2O standard, suggesting an average Cu oxidation state of +1 (Fig. 5c). Extended X-ray absorption fine structure (EXAFS) result shows that the NO-exposed Cu sample contains a mixture of Cu and Cu2O (Fig. 5d). Slight oxidation of Cu is likely due to the exposure of the sample in the air during sample handling. After a constant current density of 5 mA cm−2 was applied under CO2RR condition, the Cu catalyst was quickly reduced to metallic Cu, suggesting that a small amount of current is sufficient to fully reduce the Cu catalyst under CO2RR conditions. XAS measurements on Cu samples exposed to NO2 and N2O also exhibited similar behaviors as the NO-treated Cu sample (Supplementary Figs. 32 and 33), confirming that the Cu catalyst remains or revert to fully metallic under reaction conditions after NOx is removed from the CO2 stream.

Moreover, ex-situ SEM images were obtained at various points of the experiment to evaluate the impact of NOx on the catalyst morphology. SEM images of the spent catalysts after the exposure to NOx impurities (t = 1 h and 3 h) exhibit minimal changes in Cu and Ag catalysts (Supplementary Figs. 1, 34, and 35). Although an increase in particle size was observed in the case of Sn catalysts (Supplementary Figs. 1 and 36), the Sn sample obtained after 1 h of CO2 electrolysis in the absence of NOx also showed a similar increase in particle size (Supplementary Fig. 37). The Sn particles likely aggregated to lower the surface energy under CO2RR condition regardless of NOx, and therefore, NOx impurities are not the primary cause of the size change of the Sn particles during CO2RR. These results suggest that the presence of NOx during CO2RR has a negligible impact on the catalyst morphology.


In summary, we investigated the influence of various NOx (i.e., NO, NO2, and N2O) in CO2RR on Cu, Ag, and Sn catalysts in a flow cell. The presence of NOx impurities reduced the CO2RR FE due to the preferential reduction of NOx over CO2. The impact of NO and NO2 is more severe than that of N2O in CO2RR due to the greater number electrons involved in NORR and NO2RR compared to N2ORR. The major NORR products are NH3, NH2OH, N2, and N2O, in which the selectivity varies among different catalysts, whereas N2O is primarily reduced to N2. Despite the loss of CO2RR FE, a small amount of NOx in the CO2 feed does not alter the metallic nature of the catalyst under CO2RR conditions as demonstrated by the XPS and XAS measurements. Furthermore, although high concentrations of NOx may be detrimental to CO2RR, NOx at typical concentrations of flue gases is compatible with CO2RR, causing small losses in CO2RR FE. NOx removal process, which is a relatively mature technology, may also be employed to ensure CO2RR operation at maximum efficiency. This work not only demonstrates the effect of a trace amount of NOx impurities that are often present in the industrial CO2 point sources on the most commonly studied metal catalysts, but also offers new insights on the electrochemical reduction of NOx, which has rarely been explored in the literature.


Electrode preparation

Commercial Cu (25 nm, Sigma-Aldrich), Ag (<100 nm, 99.5%, Sigma-Aldrich), and Sn (0.1 μm, Alfa Aesar) particles were used as cathode catalysts. Commercial IrO2 (99.99%, Alfa Aesar) was used as an anode catalyst. The catalyst inks were prepared by dissolving 3 mg of the catalyst and 20 μl of Nafion (5 weight % in 50/50 water and isopropanol) in 3 mL of isopropanol. The catalyst ink was sonicated for at least 30 min, and 0.25 mg cm−2 of the catalyst was drop casted onto a Sigracet 29 BC GDL (Fuel Cell Store).

Flow cell electrolysis

The electrochemical measurements were conducted in a three-compartment flow cell with channel dimensions of 2 cm by 0.5 cm by 0.15 cm (Supplementary Fig. 2). The electrode area was 1 cm2 and the distance between the electrode and the membrane was 0.15 cm. A FAA-3-hydroxide exchange membrane (Fumatech) was used to separate electrolyte in the anode and the cathode chamber. 1 M KHCO3 was prepared by purging CO2 (Matheson, 99.999%) into potassium carbonate (99%, Alfa Aesar) and purified using a Chelex 100 sodium salt (Sigma Aldrich). After filtering Chelex 100 sodium salt, 1 M KHCO3 was used as an electrolyte for both catholyte and anolyte and was fed at 0.9 mL min−1 via peristaltic pumps (Cole Parmer). The total gas flow rate was maintained at 19.2 mL min−1 with different flow rates of CO2, Ar (Keengas, 99.999%), and NOx. For instance, 83.3% CO2 and 16.7% Ar was prepared by flowing 16 mL min−1 CO2 and 3.2 mL min−1 Ar via Brooks GF40 mass flow controllers. 83.3% CO2, 15.87% Ar, and 0.83% NOx were prepared by flowing 16 mL min−1 CO2 with 3.2 mL min−1 of 5% NO/Ar (Matheson Gas) or 3.2 mL min−1 of 5% NO2/Ar (Matheson Gas) using a 50 mL gastight syringe (1050 SL, Hamilton) via a syringe pump (New Era Pump Systems). Syringes were quickly switched to another syringe before running out of gases. Similarly, 83.3% CO2, 15.87% Ar, and 0.83% N2O were prepared by flowing 16 mL min−1 CO2, 3.04 mL min−1 Ar, and 0.16 mL min−1 N2O (99.99%, Matheson Gas). N2O was fed by using a 10 mL gastight syringe (1010 SL, Hamilton) via a syringe pump (Cole Parmer). For NO2 experiment, the gas outlet of the electrolyzer was connected to 2 M KOH (85%, Sigma-Aldrich) to scrub the remaining NO2 and additional Ar was flowed at 16 mL min−1 to carry the CO2RR products to the GC.

CV and chronopotentiometry experiments were conducted via an Autolab PG128N. For CV measurements, the electrodes were pre-reduced at 100 mA cm−2 in 83.3% CO2 and 16.7% Ar for 10 min. The half-cell potentials were measured with respect to Ag/AgCl reference electrode (Pine Research) and calculated to the RHE scale in which E (vs. RHE) = E (vs. Ag/AgCl) + 0.209 V + 0.0591 V × pH − ηIRdrop. The pH was measured at the outlet of the catholyte channel. The resistance was measured with the current-interrupt technique40, and the measured potential was manually post IR-corrected.

Product quantification

The gas products were analyzed via a multiple gas analyzer no. 5 gas chromatography system (SRI Instruments) equipped with a Molsieve 5 A and a HayeSep D column connected to a thermal conductivity detector (TCD) and a flame ionization detector (FID). Ar was used as a carrier gas with a flow rate of 19 mL min−1 and 1 mL of sample was automatically loaded to the column. The gas sample was loaded to 0.5 m HaySep D pre-column connected to 2 m Molsieve 5 A column at 0.050 min. At 0.490 min, any molecule remaining in the HaySep D precolumn was backflushed out to vent. At 2.150 min, the gas sample was automatically loaded to 2 m HaySep D column. The column temperature was maintained at 35 °C for 2.950 min, increased to 210 °C at 40 °C/min, and maintained at 210 °C until the end of the analysis. A typical GC analyses of potential CO2RR products, N2, NO, and N2O are provided in Supplementary Fig. 38. 2% H2, 1% CO, 1% CH4, 1% C2H4, 0.50% C2H6, 0.25% C3H6, 0.25% C3H8 in Ar (Matheson) was used to obtain the chromatogram of potential CO2RR products.

The liquid CO2RR products were analyzed via 1H nuclear magnetic resonance (NMR) with water suppression using a presaturation method (Bruker AVIII 600 MHz NMR spectrometer). The liquid sample was collected at the outlet of the electrolyzer and diluted to 25% in deionized water (DI). 500 μL of the diluted sample was mixed with 100 μL of 25 ppm (volume %) dimethyl sulfoxide (99.9%, Alfa Aesar), which was used as an internal standard, in D2O.

NH3 was quantified using indophenol blue method41 with UV-vis spectroscopy (Nanodrop 2000, Thermo Scientific). 100 μL of the sample was mixed with 500 μL of alkaline hypochlorite solution (A1727, Sigma-Aldrich) and 500 μL of phenol nitroprusside solution (P6994, Sigma-Aldrich). The solution was incubated in the dark at room temperature for 20 min. 2 μL of the solution was pipetted onto the pedestal, and the absorbance was measured by UV–vis spectroscopy from 190 nm to 840 nm. The absorbance of the sample was measured at 630 nm, and the absorbance measured at 830 nm was subtracted to remove the background. The calibration curves were obtained using different concentrations of ammounium hydroxide (NH4OH; 28.0–30.0%, Sigma Aldrich) in 0.25 M KHCO3 (Supplementary Fig. 13).

NH2OH was quantified using a procedure modified from a procedure reported by Afkhami et al42. with UV–vis spectroscopy (Nanodrop 2000, Thermo Scientific). Neutral red solution was prepared by dissolving 200 mg of neutral red (Sigma-Aldrich) in 100 mL DI. Iodate solution was prepared by dissolving 1.00 g of potassium iodate (KIO3, 99.995%, Sigma-Aldrich) in 100 mL DI. A total of 500 μL of sample was mixed with 250 μL of 3.0 M sulfuric acid (Fisher Scientific) and 250 μL of iodate solution. After 5 min at room temperature, 500 μL of neutral red solution was added to the solution. The solution was incubated at room temperature for 20 min. In total 2 μL of the solution was pipetted onto the pedestal, and the absorbance was measured by UV–vis spectroscopy from 190 nm to 840 nm. The absorbance of the sample was measured at 510 nm, and the absorbance measured at 800 nm was subtracted to remove the background. The change in absorbance was determined by subtracting the absorbance of the sample solution from the absorbance of the solution with 0 mg L−1 NH2OH. The calibration curves were obtained using different concentrations of hydroxylamine (50 wt % in H2O, Sigma Aldrich) in 0.25 M KHCO3 (Supplementary Fig. 14).

Flow electrochemical mass spectrometry (FEMS)

An identical flow cell with an entrance for the MS probe at the top of the gas channel was used for the FEMS measurement (Supplementary Fig. 16). The probe consisted of a PEEK capillary with inner diameter of 0.25 mm with PTFE membrane attached at the tip of the capillary. The PTFE membrane with a pore size of 200 μm was used to prevent the entry of aqueous electrolyte, while allowing gaseous and volatile products to enter the MS chamber. The distance between the probe and the cathode was kept constant. The electrodes were pre-reduced at 10 mA cm−2 for 5 min in Ar before the introduction of 0.83% NO. The products were detected by a Hiden Quadrupole mass spectrometer (MS). The mass fragments were detected by a secondary electron detection voltage of 1700 V with an ionization potential of 70 eV and emission current of 200 A. m/z of interest was tracked over the course of the experiment, in which a constant potential was applied for approximately 2 min starting at t = 1.5 min. For the deconvolution of m/z = 17 signal, m/z = 17 signal from the water was first determined using m/z = 18 signal. Next, the contribution from water to m/z = 17 signal was subtracted from the observed m/z = 17 signals to obtain the signal from ammonia. For the deconvolution of m/z = 28 and 44 signals, m/z = 28 and 44 signals from CO2 in the electrolyte was first determined using m/z = 12 signal. m/z = 12 signal was smoothed using the Savitzsky-Golay method with a window of 30 data points to reduce the oscillations in the signal prior to deconvolution. Next, the contributions from CO2 to m/z = 28 and 44 were subtracted from the observed m/z = 28 and 44 signals to obtain the signals from NORR products. m/z = 44 signal corresponded to the signal from N2O, and this was used to calculate the contribution of N2O to m/z = 28 and 30. Lastly, the contributions of N2O to m/z = 28 and 30 were subtracted from the m/z = 28 and 30 signals from NORR products, respectively, to yield N2 and NO signals, respectively. All deconvolution was conducted using MATLAB. Mass spectra of NH4OH, NO, N2O, N2, H2O, and CO2 used for the deconvolution were obtained using the same MS equipment (Supplementary Fig. 17).

Material characterization

For SEM and XPS measurements, the electrodes were first taken out of the electrolyzer after electrolysis at desired time points. The electrodes were dried in the vacuum oven (MTI Corporation) for up to three days before SEM images were acquired with Auriga 60 CrossBeam (1.5 kV). The electrodes were quickly transported to the XPS equipment (K-alpha Alpha X-ray photoelectron spectrometer system, Thermo Fisher Scientific) after drying in the vacuum oven for 5 min. The electrodes were exposed to air for less than 20 min. High-resolution XPS measurements were obtained at pass energy of 20 eV with a step size of 0.1 eV. Flood gun was turned on. Cu 2p, Ag 3d, and Sn 3d were scanned 10 times while N 1 s was scanned 30 times. Four different spots were scanned and averaged. All peaks were fitted using Thermo Avantage software with adventitious carbon referenced to the C1s peak at 284.8 eV.

XAS measurement was performed at the 8-ID Beamline of the National Synchrotron Light Source II at Brookhaven National Laboratory (BNL). The electrodes were taken out of the electrolyzer at 1 h (after exposure to NOx for 0.5 h) during a 100 mA cm−1 constant current CO2RR experiment with the introduction of NOx (Fig. 2a–c and Supplementary Figs. 6 and 7). In the case of NO2, the samples were exposed to 0.23% NO2 instead due to the availability of the gas at the time of the experiment. The electrodes were quickly stored in vials filled with Ar and the vials were tightly sealed with Parafilm at the home institution. The electrodes were transported to the Brookhaven National Laboratory (New York, USA) and were loaded into a XAS batch cell, which was fabricated from Teflon and 304 stainless steel, with a Kapton film window for high transmissivity for X-ray measurements (Supplementary Fig. 31). The electrodes were exposed to air for ~20 min before the measurement. Pt wire and Ag/AgCl were used as a counter and a reference electrode, respectively. 1 M KHCO3 was used as an electrolyte and CO2 was flowed at 10 mL min−1. XAS data were analyzed using the IFEFFIT package, which included ATHENA and ARTEMIS43.

Data availability

All data needed to evaluate the conclusions in the paper are present in the paper and/or the Supplementary Materials. Additional data related to this paper may be requested from the authors.


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This material is based upon work supported by the Department of Energy under Award DE-FE0029868. The authors also thank the National Science Foundation for financial support (Award CBET-1803200). This research used resources at the 8-ID Beamline of the National Synchrotron Light Source II, a US Department of Energy Office of Science User Facility operated by Brookhaven National Laboratory under contract no. DE-SC0012704. The authors acknowledge E. Stavitski (8-ID Beamline, NSLS-II, Brookhaven National Laboratory) for assistance in X-ray absorption spectroscopy measurements.

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F.J. supervised the whole project. B.K. and F.J. designed the experiments and collectively wrote the paper. B.K. conducted the electrochemical and XPS measurements. B.K. and W.C. prepared electrodes. B.K. and B.H. performed the FEMS study. H.S. conducted the SEM measurements. H.S., E.J., and S.O. performed the XAS experiments. All authors commented on the manuscript.

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Correspondence to Feng Jiao.

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Ko, B.H., Hasa, B., Shin, H. et al. The impact of nitrogen oxides on electrochemical carbon dioxide reduction. Nat Commun 11, 5856 (2020).

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