Spectroelectrochemical insight into copper cobalt catalysts for CO_2 and nitrite co-electroreduction to urea

Domain: physics
Published: 2026-01-17
Editorial outcome: major_revision

Abstract
Electrochemical CO and nitrite co-reduction provides a sustainable urea synthesis route but remains limited by low selectivity and an undecided C–N coupling mechanism. Here, we report co-sputtered bimetallic Cu–Co catalysts that facilitate urea formation via a tandem relay mechanism. The optimal Cu:Co ratio of 1:1 achieves a urea yield rate of 61 ± 6 mmol h⁻1 g⁻1 at –1.2 V vs. RHE under neutral pH, emphasizing the importance of proton balance in sustaining proton-coupled electron transfer. In situ synchrotron-based infra-red and Raman spectroscopy monitor the dynamic evolution of *CO, *NH , and C‒N intermediates. In situ X-ray absorption spectroscopy indicates the structural stability of metallic Cu and Co active sites. Density functional theory calculations suggest that *COOH + *NH coupling initiates urea pathway, with *NH CO formation as the potential-determining step. This study integrates rational catalyst design and in situ spectroelectrochemical analysis to advance understanding of electrochemical C–N coupling for urea synthesis.

Introduction
Urea (CO(NH ) ) is the most utilized fertilizer globally with market size around $129 billion in 2023, which it is projected to grow to around $161 billion by 2032 1. Currently, the primary method for commercial urea production is the Bosch-Meiser process, which involves the reaction of ammonia (NH ) and CO under elevated temperature (150–200°C) and pressure (100–200 bar) 2. NH is typically derived from the Haber-Bosch process, which relies on hydrogen sourced from fossil fuels, making upstream hydrogen production the main contributor to energy consumption and carbon emissions 2. As a result, the overall urea production remains carbon-intensive, generating ~910 kg of CO for every ton of urea produced 3,4.

Electrochemical co-reduction of CO and NO⁻, such as NO ⁻ or NO ⁻, offers a promising, more sustainable alternative. These NO⁻ species can be sourced from exhaust flue gas, chemical production by-product, wastewater streams, or synthesized via cold plasma oxidation of N 5–10. Such approaches could enable decentralized and potentially carbon-neutral urea production. However, urea electrosynthesis from CO and NO⁻ is a complex multistep process involving multiple electron transfer (Tables S1, S2, Supplementary Data 1) and competing with mechanistically simpler reduction processes, viz. the hydrogen evolution reaction (HER), CO reduction reactions (CO RR), and NO ⁻/NO ⁻ reduction reactions (NO ⁻RR/NO ⁻RR or NO⁻RR) 2,11. Notable challenges include limited CO adsorption capacity 12, sluggish dissociation of the C=O bond 12,13, and difficulty in balancing adsorption and activation of both carbon- and nitrogen-derived species 2,14.

At the core of these challenges lies the C–N coupling step, whose mechanism remains poorly resolved due to interplay of multiple surface-bound species and dynamic changes in catalyst structure under reaction conditions. Proposed C–N coupling pathways involve intermediates such as *CO, *NO, *NH , and *NH, with key steps often limited by proton-coupled electron transfer (PCET) barriers 11,15–17. Elucidating these pathways—and understanding how the catalyst surface evolves during operation—requires in situ spectroelectrochemical characterization techniques, such as Raman, infrared, and X-ray absorption spectroscopy (XAS), which can directly capture transient intermediates and structural changes 12,13. Achieving the most energetically favorable pathway requires precise tuning of intermediate adsorption energies, a task that may be inherently challenging for a monometallic catalyst. This has led to growing interest in multimetallic designs, where distinct active sites can preferentially stabilize different intermediates and facilitate sequential steps. Indeed, such strategy has been successfully applied in the oxygen evolution reaction (OER) 18,19 and CO RR 20,21 fields, offering a blueprint for tackling the complexity of urea electrosynthesis.

Copper has been widely studied due to its ability to facilitate CO and NO⁻ activation, moderate hydrogen binding, and tendency to form key intermediates such as *CO and *NH 22–27. However, Cu suffers from limited urea selectivity and deactivation caused by strong adsorption of nitrogen intermediates 28–32. Cobalt, on the other hand, offers high conductivity, favorable *CO binding, and the ability to stabilize *H intermediates, thereby promoting downstream protonation or hydrogenation steps 31,33–36.

These complementary characteristics led us to a working hypothesis that combining Cu and Co could couple their strengths—enhancing NO⁻ reduction to reactive nitrogen species on Cu, while stabilizing *CO and *H on Co—to accelerate C–N bond formation 31,37,38. Supporting this concept, recent work on amorphous Cu–Co boride (a-Cu CoB) demonstrated alleviation of strong N-based and *CO intermediates binding on the catalyst surface and achieved a faradaic efficiency (FE) of ca 28% with a urea yield rate of ca 310 μg h−1 mg −1 (5.2 mmol h−1 g −1) from CO and NO ⁻ 39.

In parallel, we also considered an alternative, cross-assigned tandem hypothesis—that Cu predominantly improves the CO adsorption and activation, stabilizing *CO/*COOH 40,41, whereas Co primarily promotes NO ⁻ activation and stabilizes N-intermediates 42,43. To address the ambiguity and resolve the roles of the two metals, we explored tandem Cu‒Co catalysts prepared via co-sputtering. This approach enables atomic-level mixing of two metals, providing different but complementary active sites for key reaction steps. We anticipated that such atomic-scale integration would enhance PCET, promote the formation of key intermediates, thereby optimizing the C‒N coupling process 44–46.

To test these hypotheses, we first performed systematic electrochemical studies under varying electrolyte conditions—including pH and nitrogen-source species—to evaluate urea formation and identify optimal conditions. Guided by these results, we employed in situ spectroelectrochemical techniques—in situ Raman, synchrotron radiation Fourier-transform infrared spectroscopy (SR-FTIR), and XAS—to monitor the formation and evolution of reaction intermediates directly on the catalyst surface under operating conditions. Finally, density functional theory (DFT) calculations were performed under the optimized reaction conditions to provide atomic-scale insights into the roles of Cu and Co.

Together, these results support the alternative cross-assigned mechanism, rationalizing the observed selectivity where Cu predominantly stabilizes CO -derived intermediates (*CO/*COOH), and Co drives *NO ⁻ reduction and stabilizing *NH for C‒N coupling. This integrated experimental–theoretical approach aims to unravel the C–N coupling mechanism and guide the rational design of efficient catalysts for sustainable urea electrosynthesis.

Results
We varied Cu and Co stoichiometry within the synthesized electrodes via co-sputtering metals to form CuCo , CuCo, and Cu Co onto carbon fiber paper (CFP) substrates (collectively referred to as Cu Co ) and characterized as-prepared materials using a range of physiochemical characterization methods. Inductively coupled plasma mass spectrometry (ICP-MS) analysis confirmed that metals were deposited close to expected ratios (Table S3, Supplementary Data 2).

Scanning electron microscopy (SEM) coupled to energy-dispersive X-ray spectroscopic (EDS) analysis confirms uniform distribution of the Cu and Co over the surface of carbon fibers (Fig. 1a–c and Figs. S1, S2, S4). EDS analysis under low-magnification transmission electron microscopy (TEM) conditions indicates homogenous mixing of Cu and Co within the sputtered materials with no observable segregation within detection limit of individual metallic phases (Fig. 1d, e and Figs. S3–S5). Higher resolution TEM of the as-prepared catalyst indicates the presence of well-defined crystal facets of a face-centered cubic (fcc) structure with an interlayer spacing of ~1.9 and 2.1 Å. This is consistent with the fcc (200) and fcc (111) facets, respectively, of both copper and cobalt metal phases (Fig. 1f and Figs. S6) 47. The selected area electron diffraction (SAED) patterns further validate the presence of these phases (Fig. 1g and Fig. S6). Moreover, these crystal facets remain observable even after electrochemical testing, as evidenced by the post-reaction TEM analysis (Fig. S7). X-ray diffraction (XRD) patterns of all samples exhibit a prominent peak at 43.6°, which is consistent with the (111) plane of the fcc Cu and Co metals (ICDD 00-004-0836 and 01-071-4238) (Fig. S8). A peak at 39.6° is most likely associated with the 200 signals of the monoclinic CuO phase (ICDD 01-070-6827).

X-ray photoelectron spectroscopy (XPS) was employed to elucidate the surface chemical composition of the as-prepared Cu Co catalysts (Fig. 2a, b and Figs. S9–S14), with monometallic sputtered Cu and Co serving as references. The survey spectrum reveals distinct peaks corresponding to C 1s, O 1s, Cu 2p and Co 2p, with relative intensities of Cu 2p and Co 2p peaks following the bulk composition (Fig. S9). Fitting of the high-resolution Cu 2p and Co 2p spectra of as-synthesized catalysts using the parameters reported by Biesinger et al. 48,49 and our previous reports 50,51, as a guide displays the presence of Cu(0), Cu(I), Cu(II), Co(0), and Co(II) species, indicating oxidation of surface Cu and Co metals to Cu(I), Cu(II) and Co(II) due to air exposure (Fig. 2a, b, Figs. S11–S12, Table S4, Supplementary Data 3) 52,53. High-resolution Cu LMM Auger spectra further confirm the presence of Cu(0), Cu(I), and Cu(II) species (Fig. S13). Analysis of the high-resolution O 1s spectra of Cu Co catalysts identifies lattice (529.3 eV) and non-lattice oxygen (~531.5 and 532.5 eV) (Fig. S14) 54. Post-reaction analysis of Cu Co catalyst reveals changes in surface structure after 12 h electrochemical testing (vide infra) (Fig. 2a, b and Figs. S9). In addition to Cu 2p, Co 2p, C 1s, and O 1s peaks, survey spectra detect K 2p and K 2s peaks, indicating adsorption of K+ from the electrolyte on the catalyst surface. Analysis of Co 2p, Cu 2p, and Cu LMM Auger spectra indicate further oxidation of Cu(0) and Co(0) to Cu(I), Cu(II) and Co(II) induced by the electrocatalytic testing (Figs. S11–S13) 55. Consequently, the oxidation leads to the increased peak intensity of lattice oxygen, as seen in high-resolution O 1s spectra (Fig. S14). XAS measurements of Cu Co catalyst were recorded at Cu and Co K-edges. The X-ray absorption near edge structure (XANES) and extended X-ray absorption fine structure (EXAFS) spectra (Fig. 2c–h and Figs. S15–S18) confirm that the as-prepared catalysts are dominated by metallic Cu(0) and Co(0), consistent with the XRD and HRTEM results. Fourier transformation (FT-EXAFS) and k3-weighted EXAFS spectra further support this assignment (Fig. S16). Linear combination fitting (LCF) of both XANES (Figs. S19) and EXAFS (Figs. S20) indicate that the materials being composites primarily consisting of metallic Cu and Co, without evidence of intermetallic phases. In addition to the dominant metallic phases, minor Cu O/Cu(OH) and Co O /Co(OH) species are also present as detectable oxide/hydroxide phase. After electroreduction, XANES spectra (Figs. S15, S17–S18) show slight oxidation, evidenced by an increase in edge intensity, consistent with XPS results. Nevertheless, FT-EXAFS at both Cu and Co K-edges confirms that Cu–Cu and Co–Co metallic interactions remain predominant (Fig. S17–S19), with LCF analysis indicating that the post-reaction catalysts are still primarily metallic, with fractions of oxidized Cu and Co species present (Fig. S19, S20).

Raman spectroscopy was also conducted to characterize the components of the Cu Co catalysts (Fig. S21). At 50% laser power, the as-prepared CuCo catalyst exhibits peaks indicative of Co(OH) at ~530 and 600 cm⁻1, Cu O at ~147 and 218 cm⁻1, and CuO at ~294 cm⁻1 56–60. These oxide/hydroxide features are attributed to the oxidation of metallic Cu and Co resulting from air exposure 61,62. Under power-dependent Raman acquisition, additional CuO and Co O modes emerge only at higher laser power (>50%) that may indicate a laser-induced oxidation component 63,64. Meanwhile, LCF of EXAFS spectra (Fig. S20) place an upper bound of ≤8% Co O in Cu Co catalysts, implying that any pre-existing Co O is minor near-surface fraction below bulk XAS detectability. Taken together, power-dependent Raman and LCF EXAFS spectra support that the Co O bands observed by Raman arise primarily from laser-induced oxidation, possibly superimposed on a very thin surface Co O layer.

Comprehensive characterization confirmed that as-prepared catalysts predominantly comprise metallic Cu(0) and Co(0) states. Upon electrochemical testing, partial surface oxidations to Cu(I), Cu(II) and Co(II) were observed, with the bulk metal phases largely retained, demonstrating structural stability under reaction conditions.

Electrocatalytic performance
CO and KNO were selected as the carbon and nitrogen sources for our experiments to enhance C‒N coupling based on our previous finding and the above discussions 65. The CuCo catalyst was initially evaluated for its activity towards the CO RR, NO ⁻RR, and electroreduction of CO and NO ⁻ in the CO -saturated 0.2 M KHCO solution (pH 7), Ar-saturated 0.2 M KHCO + 0.02 M KNO solution (pH 8), and CO -saturated 0.2 M KHCO + 0.02 M KNO solution (pH 7), respectively. To maintain saturation in these catholytes, CO or Ar gas was continuously injected.

Cyclic voltammetry (CV) measurements (Fig. 3a and Figs. S22a–S25) reveal the current response of CuCo catalyst under each reaction condition. Notably, distinct oxidation and reduction peaks are observed between −0.1 and 0.85 V vs. RHE in all cases (Fig. S23), indicating the formation of Cu and Co oxides/hydroxides and their subsequent reduction. Notably, CuCo exhibits quasi-reversible redox behavior, in contrast to the reversible transition of pure Cu and the irreversible features of pure Co (Fig. S24) 66–69. The irreversible nature of pure Co is likely linked to Co2+/Co3+ dissolution into the electrolyte 70,71. Together, these findings suggest that Cu–Co bimetallic occupy an intermediate regime, likely due to electronic and structural stabilization at the interface 72.

Furthermore, the CV profiles show that CO RR exhibits lower current density compared to NO ⁻RR and electroreduction of CO and NO ⁻, suggesting relatively lower activity towards CO reduction (Fig. 3a). Chronoamperometry (i-t) measurements for the CO RR, NO ⁻RR, and electroreduction of CO and NO ⁻ indicate that the CuCo catalyst achieves j of approximately −44, −80, and −64 mA cm⁻2 (−0.22, −0.41, −0.33 mA cm ⁻2) at −1.2 V vs. RHE, respectively (Fig. S28-S31). NO ⁻RR testing of the CuCo catalyst over the potential range −0.6 to −1.2 V vs. RHE showed that the highest ammonium (NH +) selectivity was achieved at −0.8 V vs. RHE with FE and yield of 96 ± 4% and 340 ± 9 mmol h−1 g −1 (0.24 ± 0.006 nmol s−1 cm ⁻2) (Fig. 3b). These values are comparable to other Cu- and Co-based catalysts in non-alkaline conditions 73,74. CO RR was also conducted in the same potentials range as NO ⁻RR (Fig. 3c). At −0.6 V vs. RHE, the reaction was highly selective to formate (HCOO⁻) with a small contribution of acetate (CH COO⁻) and methane (CH ). As the potential becomes more negative, the FE for formate progressively decreases while the amounts of other products, including methanol (CH OH⁻), grows. At −1.0 V vs. RHE, the FE for CH , CH OH, and CH COO⁻ reaches 21 ± 3%, 16 ± 4%, and 12 ± 2%, respectively. The use of even more negative potential did not improve product selectivity due to the enhanced HER.

Urea electrosynthesis was conducted via co-electroreduction of CO and NO ⁻. During 12 h experiments, urea was consistently detected at −1.0 V vs. RHE and more negative potentials using both 1H NMR spectroscopy and gas chromatography with mass-spectrometry detector (GC-MS) (Fig. 3d, e and Fig. S32). Furthermore, 15N isotope-tracing experiments confirm that the produced urea originates from the NO⁻ feed (Fig. S32 and Fig. S42). The highest performance was observed at −1.2 V vs. RHE, with a maximum FE of 11 ± 2% and yield rate of 61 ± 6 mmol h‒1 g ‒1 (8.5 ± 0.8 nmol s−1 cm−2 per geometric surface area of the electrode or 0.043 ± 0.005 nmol s−1 cm −2). Although CO RR exhibits high H formation at this potential (Fig. 3c), introducing NO ⁻ suppresses HER through two effects: (i) NO ⁻-driven electron–proton sinking and site blocking—NO ⁻ might occupy H-binding sites and lower HER 75,76. This is consistent with the additional redox couple that appears in CV curves for CuCo upon NO ⁻ addition (Figs. S23–S25), indicative of NO ⁻-metal interaction on the catalyst surface 77. (ii) Tandem perimeter kinetics of Cu‒Co that co-generate *CO/*COOH and *NH in close proximity (Fig. S33), as confirmed through in situ SR-FTIR (Fig. S40), thereby enabling high urea rates despite the H -favoring bias 75.

Time-dependent studies revealed that urea was first detectable after 5 h, with an FE and yield rate of 10 ± 1% and 56 ± 3 mmol h‒1 g ‒1, respectively (Fig. 3f). The stability test further demonstrated that the CuCo catalyst can steadily produce urea over 48 h (Fig. S30, S31). In addition to urea, C and other C‒N products, such as ethanol (C HOH), acetate, and methylamine (CH NH ) were detected during the experiments at −1.0 V vs. RHE, although at lower FE and yield rates (Fig. S30). The maximum FEs for methylamine, ethanol, and acetate were observed at −1.3 V vs. RHE, measuring at 1.2 ± 0.1%, 3.0 ± 0.7%, and 1.0 ± 0.1%, respectively (Fig. S29).

When NO ⁻ was used as a nitrogen source instead of NO ⁻, similar products were formed but at lower faradaic efficiencies and yield rates (Fig. 4a and Fig. S30). Urea was detected only at more negative potentials (≤ −1.1 V vs. RHE) and achieved maximum values of 8 ± 1% and 26 ± 4 mmol h−1 g −1 (0.018 ± 0.003 nmol s−1 cm −2) at −1.3 V vs. RHE (Fig. S31). This difference is consistent with our previous findings, where the activation and reduction of NO ⁻ require lower overpotentials compared to NO ⁻, enabling the preferential reduction of prior reactant 65. Supporting evidence was provided by in situ SR-FTIR spectra (Fig. S42), in which characteristic urea-related vibrations (*NCO, *CO carbonyl, and *CN) were observed more weakly in CO + NO ⁻ than in CO + NO ⁻ experiments 14,16,78.

To further probe nitrogen source effects, hydroxylamine (NH OH) was tested. Instead of urea, the reaction yielded ammonium carbamate ([NH ][NH COO]) with FE of 16 ± 3% and yield rate of 90 ± 5 mmol h‒1 g ‒1 (0.064 ± 0.004 nmol s−1 cm −2) at −1.2 V vs. RHE (Fig. 4a). This suggests that the dominant process is electrochemical reduction of NH OH to NH +, while the carbamate anion forms chemically through reaction of hydroxylamine with CO 79,80. The absence of ammonium carbamate in NO ⁻ and NO ⁻ experiments implies that hydroxylamine is not an intermediate in the co-reduction pathway, or if transiently formed, is rapidly converted into other species. Overall, these results establish that NO ⁻ is a more effective nitrogen source than NO ⁻ or NH OH for urea electrosynthesis. Compared with reported catalysts (Table S5, Supplementary Data 4), the CuCo system achieves competitive urea yield rates under prolonged operation, demonstrating its effectiveness as a tandem Cu–Co catalyst for selective C–N coupling.

We also investigated electroreduction of CO and NO – at −1.2 V vs. RHE using monometallic Cu and Co catalysts (Fig. S30 and Fig. S33) as control references. The sputtered Cu catalyst produced no detectable C–N products, instead yielding formate, ethanol, and propanol (Fig. S29, S47, S48). This outcome contrasts with many reports describing strong C–N coupling on Cu and underscores a more nuanced picture 15,81. In the literature, high urea selectivity is typically associated with oxide-derived Cu, where mixed Cu(0)/Cu(I) and defect-rich surfaces stabilize polar intermediates such as *CO, *NCO and *NH CO 27,82. By comparison, sputtered Cu typically generates thin, relatively smooth films with fewer defect sites, thereby limiting the stabilization of key intermediates required for C‒N coupling 83. Consequently, metallic Cu favors CO adsorption and *CO/*COOH formation 40, consistent with the results observed in the in situ SR-FTIR spectra (Fig. S42). For pure Co, the main products were ammonium and H (Fig. S33), with no evidence of C–N products, also corroborated by in situ SR-FTIR (Fig. S42).

These controls underscore that neither Cu nor Co alone is sufficient for effective C–N coupling. Instead, the tandem interface of Cu–Co is essential: Cu sites efficiently adsorb and activate CO to *CO, while Co sites accelerate the stepwise reduction of NO ⁻ to *NH /*NH . The spatial co-generation and proximity of *CO and *NH at the Cu–Co perimeter enhances the likelihood of C–N bond formation, leading to higher urea activity and selectivity than either metal alone.

Electrochemical CV measurements further highlight this synergy, as the Cu–Co bimetallic catalyst exhibits intermediate quasi-reversible behavior between pure Cu and Co (Fig. S24). This suggests that Cu suppresses Co dissolution by stabilizing its redox transitions, while Co mitigates Cu over-oxidation 78,84,85. Such mutual stabilization likely arises from electronic and structural interactions at the Cu–Co interface, further reinforcing the unique cooperative role of the tandem system. Consistently, electrochemical impedance spectroscopy (EIS) reveals that CuCo has a lower charge transfer resistance compared to pure Cu and Co (Fig. S27), underscoring its enhanced electronic synergy that facilitates CO and NO ⁻ and underpins the C–N coupling performance of the bimetallic catalyst. This effect was not reproduced by simple physical mixing: a drop-cast Cu/Co composite produced only ammonium, formate, H , CO, and C alcohols (Fig. S33), with no detectable urea. This finding demonstrates that atomic-level proximity achieved by co-sputtering is critical for cooperative interactions and the relay-type mechanism required for urea synthesis.

Given the importance of Cu and Co mixing, we also explored the effects of the Co concentration/ratio in the Cu /Co matrix on the FE and yield rate of urea. As compared to CuCo, the Cu Co composition produces lower j at potentials more negative than −1.0 V vs. RHE (Fig. 4b). Methylamine is detected as the main C‒N product, albeit at a low FE of less than 1%, within the potential range of −1.0 to −1.2 V vs. RHE (Fig. 4c). The increased amount of Cu in the catalyst favored the C‒C coupling, producing C products such as ethanol and n-propanol. At −1.2 V vs. RHE, the FEs for ethanol and n-propanol achieve 22 ± 4% and 2.1 ± 0.4%, respectively. At −1.3 V vs. RHE, the FE of both products improved to 25 ± 2% and 2.4 ± 0.7%, accordingly. Conversely, the use of the Co-rich CuCo catalyst enhances j for the electroreduction of CO and NO ⁻ at potentials more negative than −1.0 V vs. RHE as compared to CuCo (Fig. 4b). However, neither urea nor methylamine were produced. The predominant products from the CuCo catalyst are ammonium, formate, and H . These results highlight the importance of a specific Cu:Co ratio in balancing the adsorption and reduction of reactants, thus promoting urea generation without overshadowing the process with excessive side reactions (HER, CO RR, and NO ⁻RR). Moreover, all Cu Co catalysts demonstrated good stability over the course of 12 h chronoamperometric experiments for the co-reduction of CO and nitrite, as confirmed by ICP-MS analysis (Fig. S35), showing less than 10 wt% loss of Cu and Co. This loss is considered low compared to the literature 30,86.

Beyond tuning the Cu:Co ratio, we also explored the effects of the ion-exchange membrane and pH of the catholyte and anolyte on the effectiveness of the C‒N coupling for the urea generation. Control experiments were performed by adjusting the pH of both catholyte and anolyte to acidic, neutral, and alkaline conditions under continuous CO -gas flow (Table S6, Supplementary Data 5). No urea was detected after co-reduction of CO and NO ⁻ at −1.2 V vs. RHE under either acidic (pH 3 and 5, K SO , Nafion membrane) or alkaline conditions (pH 10 and 13, KOH, anion exchange membrane (AEM)). Under alkaline conditions, formate was the dominant product, acidic electrolytes suppressed C‒N coupling and steered the reaction toward H and ammonium generation (Fig. S36). Urea formation was consistently observed only when KHCO was used as buffer in both catholyte and anolyte, together with Nafion as membrane (Fig. 4d and Fig. S31). Notably, replacing Nafion membrane with AEM under identical conditions eliminated urea production (Table S6, Supplementary Data 5). These results indicate that controlled proton availability is critical for regulating surface protonation steps, consistent with a PCET pathway rather than mediation by surface-bound hydrogen (*H) 87–89. This conclusion is further supported by in situ SR-FTIR measurements (Fig. S41), which show the absence of broad vibrational signatures characteristic of *H-related species in the 3300–3400 cm−1 region 90,91.

To rationalize these trends, we consider the electrolyte-dependent mechanistic factors. Neutral KHCO maintains relatively high dissolved CO concentrations, in contrast to alkaline conditions where CO is converted to HCO ⁻/CO 2⁻, thereby enabling effective *CO/*COOH formation 92. In acidic electrolytes, *CO can form but HER dominates surface sites and mass transport, strongly inhibiting C‒N coupling 93. For NO ⁻ reduction, neutral pH provides sufficient proton activity to drive *NO hydrogenation to *NH with out excessive HER, whereas limited proton availability in alkaline electrolytes slows hydrogenation steps and reduces *NH coverage 94–96. Furthermore, neutral KHCO mitigates cathodic alkalinization, stabilizing intermediates such as *NCO and *NH CO without severe *OH blocking that occurs in strongly alkaline media 97–99.

Anolyte variation further highlighted the importance of proton flux (Table S6, Supplementary Data 5). Crucially, while the anolyte pH was varied to modulate charge transport, the catholyte bulk pH was maintained as intended, confirmed by negligible shifts (≤ ± 0.2 pH units) before and after testing. Maintaining the catholyte at neutral KHCO but lowering the anolyte to pH 5 enable detectable urea (FE 2 ± 0.2%, yield rate 14 ± 2 mmol h‒1 g ‒1; Fig. 4F and Fig. S37), whereas excessive proton concentration (pH 3) diminished CO availability and enhanced competing H /NH + pathways. This delicate balance in H⁺ concentration reflects the interplay between CO activation and NO ⁻ hydrogenation, in which optimal PCET-driven protonation enables co-coverage of *CO/*COOH and *NH 23,88,100,101. Consistent with this interpretation, in situ SR-FTIR detected their co-existence alongside urea vibrational bands under neutral conditions (Fig. S41-S42). These results indicate that urea formation depends on both an optimal Cu–Co surface and controlled proton management across both membrane and electrolyte compartments.

The role of electrolyte pH on catalyst stability is further reflected in CVs recorded in the non-catalytic potential window (Fig. S34). Acidic electrolytes promoted rapid Cu/Co redox cycling, consistent with accelerated dissolution 102,103, whereas alkaline electrolytes induced broad, hysteretic oxidation waves characteristic of irreversible oxide/hydroxide formation 99. By contrast, neutral KHCO displayed only weak and broadened transitions, indicating moderate surface dynamics. These CV signatures demonstrate that acidic and alkaline conditions drive aggressive redox activity and oxide formation, compromising structural integrity, whereas neutral conditions preserve a stable Cu–Co interface.

The instability was corroborated by ICP analysis of Cu and Co dissolution under strongly acidic and alkaline environments (Fig. S35). In acidic media (pH 3), Cu and Co dissolution achieved 8 ± 2 wt% and 17 ± 2 wt%, respectively. In alkaline media (pH 13), dissolution was lower (3 ± 1 wt% for Cu and 8 ± 2 wt% for Co). Such losses reduce the long-term stability and performance of the catalysts. Finally, we observed that urea is prone to decomposition into ammonium and formate/formic acid under highly acidic (pH 3, K SO ) and alkaline (pH 13, KOH) conditions (Fig. S38), especially at low concentrations.

In situ spectroelectrochemical studies
To correlate surface intermediates with catalyst structure, we combined in situ Raman and SR-FTIR with XAS. In situ Raman spectroscopy (900–3400 cm−1) tracked the evolution of intermediates during CO and NO ⁻ electroreduction (Fig. 5a and Fig. S39). Under open circuit potential (OCP), bands from bidentate nitrite (νNO , ~1360 cm−1) and buffer solution (HCO ⁻, ~1062 cm−1) were observed and persisted at −1.2 V vs. RHE 104–106. After 20 min at −1.2 V vs. RHE, new features at ~730 (NO bending) and ~2116 cm⁻1 (CO stretching) appeared, consistent with the reduction of NO ⁻ and CO , respectively 107–109. Urea related vibrations at 1002 cm−1 (C‒N) and 575 cm−1 (amide OCN) emerged after 30 min, but were absent in the control CO RR or NO ⁻RR (Fig. S38), indicating its specific association with C‒N coupling 107. Similar intermediates were identified under different applied potentials (Fig. S39), with C‒N bond and OCN bending identified from −0.9 V vs. RHE. Nevertheless, the amine group intermediate signals (e.g., *NH ) resulted from further reduction of NO ⁻ and NO are likely undetectable in Raman spectroscopy due to weaker NH vibrations and low laser power 110.

In situ SR-FTIR analysis (900–3400 cm−1) provided complementary evidence of intermediate formation at various applied potentials (Fig. 5b, c). Bands assigned to *NH bending mode (~3115 cm−1), *CO stretching (~2065 cm−1), *NH rocking vibration (~1150 cm−1) and *NH rocking/wagging mode (~930 cm−1) appeared from −0.8 V vs. RHE, indicative of CO and NO ⁻ reduction 14,82,111–113. An amide-related *NCO vibrations (e.g., *NHCO, *NH CO, or *NH COOH), emerged around 1694 cm⁻1 from −0.9 V vs. RHE, suggesting the onset of C‒N coupling formation 14,16,78. Vibrational features assigned to the carbonyl (*CO) and *CN bonds of urea were observed at ~1590 cm−1 and 1455 cm−1, respectively, over the potential range of −0.9 V to −1.3 V vs. RHE, align closely with the C=O and C‒N stretching frequencies of the urea reference (1589 cm−1 and 1457 cm−1, respectively; Fig. S40). In contrast, control CO RR and NO ⁻RR experiments display only *CO/*COOH and *NH /*NH features (Fig. S41), without any C‒N features, thereby indicating that *CO/*COOH and *NH likely act as precursors for C‒N coupling 14,16,17,78.

Catalyst comparison highlighted the role of the Cu‒Co interface. Urea-associated carbonyl and *CN vibrational features were exclusively observed for CuCo (Fig. S42). Pure Cu exhibited mainly *CO (~2065 cm−1) and *COOH (~1720 cm−1) bands associated with gaseous CO and formate formation, along with a C‒O stretching band (*CO) at ~1030 cm−1 corresponding to methanol 114. These results indicate that CO RR dominates on pure Cu, consistent with earlier observations. In contrast, pure Co displayed no carbonyl or *CN bands, yielding mainly formate and ammonia. These results indicate that the Cu–Co interface promotes the C‒N coupling steps leading to urea.

Isotope labeling with 15NO⁻ confirmed the nitrogen origin of the *CN band (Fig. S43). Spectra for 14NO⁻ and 14NO⁻ were nearly identical, apart from an additional *NO⁻ feature at ~1340 cm−1 113. Substitution with 15NO⁻ shifted nitrogen-containing vibrational modes to lower wavenumbers, with the *C15N feature of urea detected at 1411 cm⁻1, while the carbonyl peak remained unchanged 115. This isotopic shift confirms that the *CN signal originates from NO⁻-derived nitrogen.

Catalyst structure and dynamics were probed by in situ Raman and XAS. In situ Raman spectra in the 300–1000 cm−1 region showed no oxides or hydroxides signals under negative potentials (Fig. S39), suggesting that the active sites during the electroreduction are in the metallic forms of Cu and Co. Only at OCP, surface hydroxides (Co(OH) and Cu(OH) ) were detected at ~502.5 cm−1, implying the interaction between Cu/Co and the electrolyte 31,58,116.

The results from in situ Raman are corroborated by in situ XAS analysis at Cu and Co K-edges. The in situ XANES of CuCo catalyst at Cu K-edge (Fig. 5d and Fig. S44) reveals a characteristic XANES consistent with Cu(0) and a 3d10 electron configuration 25,117. A slight shift to lower effective nuclear charge in the rising edge between 8980 and 8990 eV was observed under increasingly applied potentials, suggesting some surface copper oxides is converted to Cu(0) 118. Additionally, the main peak (white line) during in situ conditions is shifted to lower energies (Fig. S44), consistent with a transient reduction associated with adsorbed reactants during electrocatalytic reactions 119. The k3-weighted EXAFS (Fig. 5e) and normalized FT-EXAFS (Fig. 5f) at Cu K-edge confirm that the majority Cu retains characteristics of metallic Cu under negative potentials, indicating that Cu returns to its metallic state. OCP measurements before and after reaction show that CuCo exhibits Cu–O and Cu–Cu interactions at apparent distances of ~1.4 and ~2.2 Å, respectively. Additionally, LCF of XANES and EXAFS spectra (Fig. S49–S50) confirm that Cu(0) is the dominant component under cathodic potentials, with copper oxides and hydroxides (i.e., Cu(OH) ) each accounting for <10%. These minor oxides/hydroxides species re-emerge once the potential is removed, driven by re-oxidation upon contact with the electrolyte 65,120. Similar interaction patterns are also observed during the CO RR and NO ⁻RR using CuCo and Cu (Fig. S45-S51).

In situ XAS measurements at the Co K-edge, examining both near-edge and extended-edge structures, reveal a gradual decline in the oxidation state of the catalyst to metallic Co with increasing cathodic potential (Fig. 5g–i and Fig. S52). This change is marked by a subtle shift towards a lower effective nuclear charge between 7710 and 7718 eV, accompanied by a decrease in white-line intensity, both indicative of a reduction in the oxidation state of Co (Fig. 5g and Fig. S52). Unlike copper, which tends to form oxides under OCP conditions, the k3-weighted EXAFS (Fig. 5H) and normalized FT-EXAFS (Fig. 5i) at Co K-edge demonstrate that Co predominantly retains the properties of metallic Co under both OCP conditions and negative potentials. This is likely attributed to the dissolution of Co2⁺/Co3+ species into the catholyte, rendering them undetectable by XAS. The FT-EXAFS analysis further shows a consistent increase in the intensity of Co–Co coordination at an apparent distance of 2.1 Å with the application of negative potential (Fig. S52). Comparable interaction patterns are also observed for both CO RR and NO ⁻RR using CuCo and Co (Fig. S53–S55). LCF of the XANES and EXAFS spectra further verifies that metallic Co is the principal phase under both OCP and negative potential, with trace of Co O and Co(OH) during cathodic potential (Fig. S56–S58).

Together, these results indicate a strong structure-activity correlation. Minor copper oxides/hydroxides detected by LCF EXAFS/XANES may, in principle, stabilize polar intermediates such as *NCO 121–123. Similarly, cobalt oxides/hydroxides have been reported to enhance NO ⁻ activation and hydrogenation 31,124. However, these species are present only in trace amounts, and their limited surface coverage or transient lifetime appears insufficient under electroreduction conditions, as depicted in CV measurements (Fig. S23). This is consistent with the in situ SR-FTIR spectra of pure Cu, which show only weak *CN features (Fig. S42). It should also be noted that LCF provides bulk-averaged phase fractions and cannot resolve whether these minority species are surface-localized, catalytically accessible, or stable under bias; hence, their role is considered indicative rather than definitive for the active surface. Overall, the evidence indicates that Cu alone enhances CO adsorption and stabilizes *CO/*COOH intermediates but lacks sufficient *NH coverage; Co alone reduces NO ⁻ to *NH /*NH but provides inadequate *CO (Fig. S42). In contrast, at the Cu–Co interface, CO -derived *CO/*COOH from Cu and NO ⁻-derived *NH from Co are generated in close proximity, which enhances the probability of C–N bond formation. Based on these mechanistic observations, we then performed DFT modeling of the reaction pathways.

DFT calculations
Experimental results, including ex situ XAS and XPS, alongside in situ measurements of Raman and XAS, have indicated both Cu and Co metals as the dominant phases. DFT calculations are subsequently used to further understand the function of Cu and Co in the catalyst. Both Cu and Co stabilize *H intermediates on their surfaces (ΔG = −0.54 and −0.62 eV) (Fig. 6a), which suppresses H evolution but might not necessarily lead to significant *H-mediated hydrogenation. Rather, this stabilization maintains the catalyst surface in a state conducive to PCET, whereby proton transfer from the electrolyte, coupled with electron transfer, drives the hydrogenation of *CO/*COOH and *NO ⁻-derived intermediates 125. This interpretation aligns with our experimental observations that urea formation is sensitive to proton flux across the membrane, consistent with a PCET-controlled mechanism.

DFT calculations indicate that Cu and Co exhibit distinct but complementary functions within the tandem system: (i) Cu enhances CO adsorption, while Co facilitates *COOH and *CO formation during CO RR (Fig. 6b); and (ii) Co promotes NO ⁻ adsorption, whereas Cu favors *NH generation during NO ⁻RR (Fig. 6c and Fig. S59). This synergistic interaction aligns with in situ spectroscopic observation of *CO/*COOH and *NH co-appearance under proton-accessible conditions. Calculations of C–N coupling pathways (Fig. 6d) further indicate that *NH + *COOH coupling is the most thermodynamically favorable route toward urea, with the formation of *NH CO identified as the potential-determining step (ΔG = 0.78, 0.22, and 0.57 eV on Cu, Co, and Cu–Co interface, respectively; Fig. 6e). Once *NH CO is formed, subsequent PCET steps toward *NH CONH and *NH CONH proceed spontaneously (Fig. 6f).

We acknowledge that our present models likely underestimate the contribution of defect-rich or amorphous sites, which could further lower activation barriers and stabilize key intermediates 126. Future studies employing large-scale simulations or machine-learning-assisted approaches will be needed to better capture the structural complexity of sputtered catalysts 127. It should also be noted that the DFT calculations are based on idealized surface models under vacuum conditions and do not fully capture solvent effects, dynamic surface reconstructions, or potential-dependent processes that occur during catalysis.

Collectively, the DFT results are qualitatively consistent with the experimental observations, suggesting that the Cu–Co interface likely balances C–N intermediate adsorption, lowers the barrier for *NH CO formation, and facilitates proton-driven hydrogenation steps. This proposed synergy offers a plausible explanation for the experimentally observed dependence of urea formation on proton flux and electrolyte environment. Furthermore, the calculations imply that Co may mitigate the overly strong adsorption of C–N intermediates on pure Cu, while the spatial proximity of Cu and Co sites promotes the overall catalytic activity.

Methods
Materials
Catalyst precursors used to produce the co-sputtered catalyst were Cu (AJA International USA, ≥99.99%) and Co (AJA International USA, ≥99.9%). Carbon-based electrodes were using carbon fiber papers (Fuelcellstore, 5% PTFE, 0.28 mm thickness, 0.44 g cm−3 bulk density, and 78% porosity).

Electrolyte solutions were prepared using KHCO (Sigma-Aldrich, ≥99.7%), KNO (Sigma-Aldrich, ≥99.6%), KOH (Sigma-Aldrich, ≥90%, flakes), K SO (Sigma-Aldrich, ≥90%, powder), H SO (Sigma-Aldrich, 95–98%) and deionized water (Satorius™ Arium®, 18.2 MΩ cm−1 at 23±2°C). The gases for electrochemical testing used CO (Coregas, grade 2.5) and Ar (Coregas, grade 5).

Ammonium identification used Berthelot method employing NH Cl (Sigma-Aldrich, ≥99.99%), NaOH (Sigma-Aldrich, ≥98%), salicylic acid (Sigma-Aldrich, ≥99%), sodium citrates (Sigma-Aldrich, ≥99%), sodium hypochlorite solution (Sigma-Aldrich, 4.00–4.99%), and sodium nitroferricyanide solution (Sigma-Aldrich, ≥99%). The 1H-NMR for CO RR liquid product and urea identification used DMSO-d 6 (Cambridge Isotope Laboratories, D-99.9%) as a solvent and sodium acetate (Sigma-Aldrich, ≥99%) as an external standard.

Identification of urea via GC-MS was employed using chemicals including acetonitrile (Sigma Aldrich, 99.8%), N-methyl-N-(t-butyldimethylsilyl)trifluoroacetamide (MtBSTFA) (Sigma Aldrich, >97%), and toluene (Sigma Aldrich, ≥99.8%). The urea identification via NMR isotope tracing was applied using K15NO3 (Sigma Aldrich, 10 atom% 15N).

Cu O (Sigma Aldrich, ≥99.999%, anhydrous), CuO (Sigma Aldrich, 99.99%), Cu(OH) (Sigma Aldrich, ≥95%), CoO (Sigma Aldrich, ≥99.5%), Co(OH) (Sigma Aldrich, ≥95%), and Co O (Sigma Aldrich, ≥99.5%) were applied as reference materials for XAS characterization of the catalyst.

Catalyst preparation
Cobalt and copper were deposited onto the carbon fiber paper substrate by simultaneous co-sputtering DC magnetron guns (1000 W) in 1200 s to obtain Cu‒Co samples. Argon was used during the deposition in the sputter coater system (AJA Intl. – ATC-2400V). The stoichiometric ratios of Co and Cu were controlled through tuning the power of gun sources. The catalyst loading is 0.2 ± 0.02 mg cm⁻2.

Electrochemical measurements
Electrosynthesis was performed in a three-electrode customized H-cell system using a Nafion 117 membrane to separate the compartments within a gas-tight cell. A platinum (Pt) plate and saturated calomel electrode (SCE) served as the counter and reference electrode, respectively. The reference electrode was calibrated against the reversible hydrogen electrode (RHE) in the same electrolyte prior to all measurements. The calibration was performed using a clean Pt plate as the working and counter electrodes in a H -saturated electrolyte. The open-circuit potential was recorded until stabilization. This calibration procedure was repeated regularly to ensure consistency across measurements. The electrolyte consisted of Ar-saturated 0.2 M KHCO , Ar-saturated 0.2 M KHCO + 0.02 M KNO , and CO -saturated 0.2 M KHCO were utilized as the catholyte for HER, NO ⁻RR and CO RR, respectively. The catholyte for C−N coupling electrosynthesis used CO -saturated 0.2 M KHCO + 0.02 M KNO solution. The anolyte for all electrochemical measurements used 0.2 M KHCO solution but without gas saturation. All electrolytes were prepared freshly on the day of electrolysis to ensure solution stability and reproducibility.

Before experiments, the cathodic compartment of the cell was purged with the relevant gas for 10 min, achieving pH levels of 8 ± 0.2, 7 ± 0.2, and 7 ± 0.2 for NO ⁻RR, CO RR and C–N coupling electrosynthesis, respectively. The Ar or CO gas flow rate was maintained at 10 mL min−1 (9.4 standard liter per minute (SLPM)) during the HER, NO ⁻RR, CO RR and C–N coupling reactions. The gas flow rate was quantified using a flowmeter (Dwyer RM Rate-Master) integrated into the gas inlet line of the cell. Electrochemical testing was conducted with an Autolab M204 (Metrohm Autolab) electrochemical workstation. The tests for NO ⁻RR and CO RR are conducted over a duration of 1 h, while electroreduction of CO and NO ⁻ is carried out between 5 and 12 h. Fixed potential electrolysis was then carried out in triplicate. Following electrolysis, all potential were manually iR corrected, where i is the measured current and R is the electrolyte resistance (Fig. S27). The corrected potentials were then converted to the RHE reference for benchmarking with the literature using the following equation: E (V) = E (V) + 0.245 + 0.059 × pH for acidic/neutral conditions or E (V) = E (V) + 0.098 + 0.059 × pH for alkaline conditions. During the adjustment of the catholyte and anolyte to acidic and alkaline conditions, as mentioned in the main text, a 0.2 M K SO solution served as the buffer for the acidic electrolyte, with the pH being adjusted using H SO . Conversely, the alkaline electrolyte utilized a KOH solution, the concentration of which varied to achieve the desired pH levels. A Hg/HgO electrode was employed as the reference electrode in alkaline electrolytes.

The cyclic voltammetry (CV) for the electrochemically active surface area (ECSA) calculations is measured with certain potential ranges (referred to Fig. S26) at scan rate of 5, 10, 20, 40, 60, 80, and 100 mV s−1. The difference between the cathodic and anodic sweeps in the CV curves plotted versus the scan rate of the CV, then fit linearly to obtain the slope that represents the double-layer capacitance (C ). The ECSA calculation follows the equation: ECSA = C /C, where C is shown in (Fig. S26), and C s is the capacitance of an atomically planar surface (C is estimated to be 40 μF cm−2) 128,129. The ECSA for each electrode size was then calculated by the following equation: ECSA = (cm2 per cm2 ) × (cm2 ). The electrochemical impedance spectroscopy (EIS) was measured under −1.2 V vs. RHE in a CO -saturated 0.2 M KHCO and 0.02 M KNO solution with the frequency from 100 kHz to 0.1 Hz. The EIS fitting was conducted using The ZFit function in ECLab software version 11.42.

Product quantification
A Shimadzu gas chromatograph (Model 2010 Plus) equipped with both a thermal conductivity detector (TCD) and a flame ionization detector (FID) was employed to detect gas products such as H , CO, and CH , with calibration data provided in Fig. S60. For the quantification of liquid products from CO RR, 1H NMR spectroscopy was used. Specifically, ~400 µL of catholyte was combined with 200 µL of DMSO-d and analyzed using a 600 MHz 1H 1D liquid NMR spectrometer equipped with a cryoprobe over 128 scans. An external standard method utilizing 10 mmol mL−1 of sodium acetate facilitated quantification.

Ammonium levels were measured using ~0.5 mL of catholyte, employing the Berthelot reaction. The reaction involved mixing the catholyte with reagent 1 (comprising 0.4 mL of 1 M sodium hydroxide, 5 wt% salicylic acid, and 5 wt% sodium citrates). This mixture was then combined with 0.1 mL of reagent 2 (0.05 M sodium hypochlorite solution) and 30 µL of reagent 3 (1 wt% sodium nitroferricyanide solution). Following sonication and a 2-h incubation in the dark at room temperature, the ammonium concentration was quantified using a Shimadzu UV-3600 UV-vis-NIR spectrophotometer and 1 cm quartz cuvettes, measuring absorbance between 550 and 800 nm wavelengths. Calibration curve data (Fig. S61) from electrolyte solutions with known NH Cl concentrations were used for this purpose.

Furthermore, identification and quantification of urea and other liquid products in C–N coupling electrosynthesis were performed using the 1H NMR method as described by Huang et al. 131. Approximately 400 µL of liquid sample was mixed with 200 µL of DMSO-d 6 solution (D-99.9%, Cambridge Isotope Laboratories) and analyzed in a 600 MHz 1H 1D liquid NMR spectrometer (Bruker Advance) equipped with a cryoprobe over 512 scans 131. The 1H spectrum was recorded using the water suppression excitation sculpting technique, and quantification of urea was determined using the calibration curve of various urea concentration detected in the 1H-NMR method (Fig. S61). For isotope tracing of urea using the 1H NMR method, K15NO3 was employed as nitrogen sources.

Other than NMR technique, GC-MS is also used to identify urea. Approximately 100 µL liquid sample was dried using nitrogen evaporator needle system (~10 min per 50 µL). After drying, sample was derivatized to its t-butyldimethylsilyl ester form at 60°C for 1 h by adding 10 µL acetonitrile and 40 µL N-methyl-N-(t-butyldimethylsilyl)trifluoroacetamide (MtBSTFA) that consisted of 1% t-butyldimethylchlorosilane (tBDMCS). Derivatives were reconstituted in 40 µL toluene, then injected (1 µL splitless = 0.25–1.25 µL media equivalent on column) into GC-MS for analysis. If sample consists of urea, it will contain urea derivatives, Urea, N, N’-bis(tert-butyldimethylsilyl).

The faradaic efficiency (FE) and yield rate of liquid products were calculated to evaluate the catalytic performance using the following equations:

FE (%) = n × F × V aq × C Product × 100% / Q

where n is number of electron transfers, F is the Faraday constant (96485.33 C mol–1), V is the volume of catholyte solution (L), C Product is the concentration of product (M), and Q is passed charge (C).

Yield rate (mmol h −1 g −1) = C Product × V aq × MW / t × A × Γ

where MW is the molecular weight of product, t is duration of the electrolysis (h), A is the geometric surface area of the electrode (cm2), and Γ is the catalyst loading per geometric surface area of the electrode (g cm–2); C Product for this calculation is expressed in mM.

The faradaic efficiency of the gaseous products was calculated using the following equation:

FE (%) = n × F × V gas × x × 100% / Q

where V is the volume of headspace (L), and x is the concentration of the product in the headspace (mol L –1).

Material characterization
XRD patterns were obtained using a D2 Phaser Bruker instrument equipped with Cu Kα radiation (λ = 1.54 Å) across a scan range of 10° to 80° at a rate of 1.17 degrees per minute. The catalyst on CFP was positioned on a zero-diffraction plate made of silicon crystal. XPS was performed using a Thermo ESCALAB 250Xi spectrometer, utilizing a monochromated Al Kα source (energy 1486.68 eV), operated at 120 W (13.8 kV × 8.7 mA), with a hemispherical analyzer in fixed analyzer transmission mode. The vacuum chamber maintained a pressure typically between 10−9 and 10−8 mbar. Survey spectra were recorded at a pass energy of 100 eV, and high-resolution spectra for individual peaks were acquired at 20 eV pass energy. The photoelectron take-off angle was set at 90°, with a spot size of ~500 µm. Data analysis was facilitated using Thermo Scientific Avantage software. Elemental identification and quantification were achieved through survey spectra, with atomic concentrations determined using integral peak intensities and sensitivity factors provided by the manufacturer. Binding energies were calibrated to the C 1s peak at 284.8 eV, representing aliphatic hydrocarbons. The accuracy of quantitative XPS is estimated at 10–15%, with a precision often exceeding 5%. The peak in the XPS spectrum was deconvoluted using the Casa XPS software.

Morphological analysis was performed using a Zeiss Ultraplus field-emission scanning electron microscope (FESEM) operating at 3 kV and a JEOL 2100F high-resolution TEM at 200 kV. For SEM, electrode samples were mounted on specimen stubs using carbon tape. TEM samples were prepared by scraping material off the electrode surface and dispersing it onto holey carbon-coated nickel grids. Elemental mapping was conducted using a JEOL SDD energy-dispersive X-ray spectrometer coupled with the JEOL 2100F. The HRTEM and SAED analysis were performed with a JEOL JEM-F200. Sample was place on in-situ holders. This instrument achieves a lattice resolution of 0.16 nm in annular dark-field scanning STEM (ADF-STEM) mode and 0.1 nm in TEM mode at 200 kV. Additionally, the F200 features a JEOL 100 mm2 energy-dispersive X-ray (EDX) detector that offers an energy resolution of 130 eV (Mn K) and a collection angle of 0.9 steradian. Ex situ Raman measurements were performed using an inVia Raman spectrometer (Horiba JOBIN YVON US/HR800 UV) with a 633 nm laser wavelength. X-ray absorption spectroscopy (XAS) of the Cu K-edge and Co K-edge was carried out at the Australian Synchrotron (ANSTO), in the XAS and Medium Energy X-ray Absorption (MEX-1) spectroscopy beamlines. Fluorescence detection was used for powder data, with Cu foil, Cu O, Cu(OH) , CuO, Co foil, CoO, Co(OH) and Co O serving as reference samples. XAS data were analyzed using Athena software and the fitting using linear combination function. MS Excel and Origin were used for the final presentations of data and fine calibrations.

For post-reaction characterization, the catalysts were subjected to 12 h of electroreduction in CO -saturated 0.2 M KHCO + 0.02 M KNO solution. After electrolysis, the samples were collected, rinsed with deionized water, and stored in a vacuum chamber until further analysis.

In situ measurements
In situ SR-FTIR measurements were conducted at the Australian Synchrotron (ANSTO), Melbourne using an SR-FTIR setup that incorporated a ZnSe crystal as the infrared transmission window (cut-off energy 625 cm–1). This setup was paired with an infrared microscope (Bruker Hyperion 3000 with a 20x objective) and an FTIR spectrometer (Bruker 70v/s) featuring a KBr beam splitter and a liquid-nitrogen-cooled mercury cadmium telluride detector. The catalyst electrode was positioned closely against the ZnSe crystal window with a micrometer-scale gap to reduce infrared light loss. The system was operated in reflection mode with vertical infrared light incidence to maintain the integrity of the SR-FTIR spectra. Each infrared absorption spectrum was derived by averaging 256 scans at a resolution of 2 cm–1. Before conducting systematic measurements, a background spectrum of the catalyst electrode was recorded at open-circuit voltage. The potential settings for the electroreduction reaction were established at −0.8, −0.9, −1.0, −1.1, −1.2, and −1.3 V vs. RHE using CO -saturated 0.2 M KHCO and 0.02 M KNO as catholyte. Fig. S62 illustrates the in situ electrochemical cell and setup configuration.

In situ Raman measurements were carried out using an inVia Raman spectrometer system (Horiba JOBIN YVON US/HR800 UV) with a spectro-electrochemical Raman cell. The Raman cell was operated using a 633 nm argon ion laser and was equipped with an Ag/AgCl reference electrode and a Pt wire counter electrode. The cell contained an electrolyte solution of CO -saturated 0.2 M KHCO + 0.02 M KNO and was connected to an Autolab potentiostat (PGSTAT204) for simulating electrocatalytic C–N coupling.

In situ XAS measurements was carried out based on previous experiment 132. In situ Cu K-edge and Co K-edge XAS data were gathered on the multipole wiggler XAS beamline (12-ID) at the ANSTO, Melbourne. The beamline operated with an electron beam energy of 3.0 GeV and a beam current of 200 mA (top-up mode), using “mode 1” and “mode 2” for Co K-edge and Cu K-edge, respectively, with a Si(111) monochromator and focusing optics, achieving a beam flux of ~2 × 10¹¹ photons per second and a beam size of 0.5 mm × 0.3 mm. Data collection was conducted in fluorescence mode with a solid-state 18-element detector. The in situ cell (Fig. S63) was mounted on the beamline sample table, positioning the electrode at a 45° angle to the fluorescence detector. An OKEN ionization chamber measured the incident radiation, and all measurements were taken at ambient temperature (23 ± 1°C).

To minimize air absorption of fluorescent X-rays, the sample environment was maintained in a helium atmosphere by connecting a flight tube between the beamline and the cell and applying a helium overpressure between the cell and detector. Each X-ray absorption near edge structure (XANES) and extended X-ray absorption fine structure (EXAFS, k=12) spectrum required ~15 and 30 min, respectively. Data was collected at multiple sample spots to confirm homogeneity. The raw data was processed in Athena.

DFT calculations
First-principles density functional theory (DFT) calculations were performed using the Vienna Ab initio Simulation Package (VASP) 133,134. The electronic wave functions were treated using the Projector Augmented Wave (PAW) approach 135,136. The exchange correlation functionals were approximated by the generalized gradient approximation (GGA) as parameterized by the Perdew-Burke-Ernzerhof (PBE) 137. The plane-wave cut-off energy of 500 eV and 3 × 3 × 1 Gamma k-point were used for the electronic self-consistent calculation, until it was converged to 1 × 10−5 eV. The ionic relaxation steps were performed using the conjugate-gradient method (IBRION=2) and continued until the total force on each atom dropped below a tolerance of –0.05 eV/Å.

Surface investigations were studied using slab methods, for the hybrid Co (111) and Cu (111) surfaces. A vacuum region with the thickness of 20 Å was introduced perpendicular to the surface (in the z-direction). To model the bulk-surface interface, atoms in the two outermost layers were allowed to relax, while the coordinates of atoms in other layers were fixed to their bulk positions. The free energy diagrams were developed by taking into account the zero-point energy and entropy contributions. Additionally, long-range van der Waals interactions were considered using DFT D-3 method 138.

Data availability
All data supporting this study are available within the article and the Supplementary Information. Source and supplementary data are provided with this paper. All other relevant source data is available from the corresponding authors upon request. Source data are provided with this paper.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "The authors developed a CuCo bimetallic catalyst for electrochemical synthesis of urea via the co-reduction of CO2 and nitrite (NO2-). The catalyst was comprehensively characterized using a series of microscopic and spectroscopic approaches. In particular, in situ X-ray absorption spectroscopy (XAS), Raman, and infrared spectroscopy were employed to probe the dynamic chemical state and adsorbed species under reaction conditions. These experimental results, combined with theoretical calculations, provide insights into the reaction mechanism. Several control experiments further confirm that a Cu:Co ratio of 1:1 delivers the best performance in terms of urea Faradaic efficiency and yield. The authors also explored the pH effect, highlighting the role of proton balance in modulating surface *H coverage and enhancing proton-coupled electron transfer. Overall, this is a systematic study for the newly developed catalyst material for the complicated electrochemical co-reduction of CO2 and NO2-. However, several key aspects regarding the nature of the active sites and the interaction between Cu and Co remain insufficiently explored. Moreover, some data were not properly explained and clarified. Specific comments are listed as below:\n1. The Introduction Part is too lengthy, making the readers not clearly understand the core message of the manuscript. Some contents are related to noble metal catalysts, which are irrelevant to this study as the manuscript is focusing on non-noble metal catalysts.\n2. The manuscript states negligible urea activity for pure Cu, which contradicts several recent studies showing urea production from CO2 and nitrate/nitrite using Cu or Cu2O (Nat. Commun. 2025, 16, 3271; Applied Catalysis B: Environ. 2024, 340, 123265). The authors should provide more detailed discussion or justification regarding to this issue.\n3. It is interesting that only Cu and Co together with a ratio or 1:1 works well for urea production. However, the detailed interaction between Cu and Co was not well demonstrated in the manuscript. In addition, the linear combination analysis of XAS data shows the catalyst is not fully reduced under negative applied potential, with the co-existence of oxidized Cu or Co species. Their amount is also varied at different reaction conditions. Therefore, the oxidized species might play important role for catalysis, which should also be further explored.\n4. Cu 2p XPS is not sufficient to clearly distinguish Cu(0) and Cu(I). Therefore, the conclusion of existing of Cu(0), Cu(II), but not Cu(I), is questionable. The peak deconvolution of CuLMM Auger spectra in Fig. S10e is also not accurate. Please check the relevant literatures (e.g. Surf. Interface Anal. 2017, 49, 1325) for more accurate XPS data fitting.\n5. The authors conducted a lot of post-reaction characterization, including XAS, XPS and Raman. However, the detailed information of the reaction condition was not clearly provided. The phenomena of further oxidation of the catalysts compared to the pre-catalysts after electrochemical reduction conditions are strange. It is highly probable that the catalyst materials were extensively oxidized in air after electrochemical tests. Given that the authors performed in situ experiments in the following section, the post-reaction characterization seems not necessary and it may mislead the readers regarding the real active species.\n6. From Line 195 to 216, this section contains repetitive statements and should be condensed for clarity.\n7. Line 222-224, the authors attributed detected Co3O4 in Raman (but not observed in XAS) to laser-induced oxidation. However, this statement might be incorrect, as XAS is more bulk sensitive and may not detect near surface Co3O4, particularly in small quantities. The LCA data in Fig. S20 also shows the presence of Co3O4 for the other control CuCo samples, which may have higher amount of Co3O4.\n8. Regarding Raman and infrared spectroscopy data, in addition to *CO, peaks between 2000–2100 cm⁻1 can also correspond to *NO (see ACS Energy Lett. 2022, 7, 1187). This is because CO and NO are chemically similar. Isotope labeling is strongly recommended to accurately assign and distinguish these species.\n9. 15N isotope labelling experiment is very important for electrochemical synthesis of urea, as the activity of urea production might be over-estimated due to the influence of the contaminated species. Although in Fig. S28b the authors showed the 15N labelling NMR data, it would be better to quantify it to demonstrate that urea is indeed generated from nitrite.\n10. In Fig. S22, the authors attributed the change of the redox peaks under different electrolyte to interaction between the reactants to Cu/Co. However, at such potential region no catalytic adsorption and reaction were observed. Therefore, more interpretations are required for the difference of redox peaks.\n11. In Fig. S23, the redox peaks of Co seem not reversible as only reductive peaks were observed. On the contrary, both Cu and CuCo show reversible or quasi-reversible redox peaks. The related explanations were absent in the current manuscript.\n12. Figs. S24 and S25 are described as scan rate-dependent, but the actual data was shown at different cycling numbers instead. Please correct the figure labels and descriptions accordingly.\n13. The caption of Fig. S31a indicated that no NO2- was added, but why NH3/NH4+ was still generated in this condition?\n14. Linear combination analysis is missing for pure Cu in Fig. S41. Please include this to allow comparison with CuCo samples.", "reviewer_2": "The author investigated the mechanism and reaction performance of the carbon-nitrogen coupling process by changing the metal ratio, conducting in-situ characterization, and performing theoretical calculations. However, the discussion in the article is still not profound enough. I reject to accept this work in this version for publication. The main points are as follows：\n1、 At -1.2 V，the HER are very serious for sole CO2RR (fig 3C), while the CuCo catalyst show the maximum urea performance in this point. Some reasonable explanations need to be provided.\n2、 In-situ characterization was employed multiple times, but only the changes in the oxidation state of copper were analyzed, without discussing the key steps of the carbon-nitrogen coupling process from the perspective of the dynamic changes of the intermediates.\n3、 The explanation for choosing NO2- as the nitrogen source is too simplistic. Why the commonly used nitrate is ineffective or inefficient in this case requires more experimental evidence to support it.\n4、 Has the 1002 cm⁻1 signal attributed to urea in in-situ Raman been verified by control experiments? Are interferences from other nitrogen-containing species excluded?\n5、 Whether the hydrogenation process is mediated by active hydrogen or the PCET process is not clearly indicated in the article's evidence.\n6、 Overall, the performance of urea is not outstanding, and the understanding of its mechanism still needs to be further enhanced.", "reviewer_3": "This manuscript presents a compelling and thorough investigation into the co-electroreduction of CO2 and NO2– to urea using co-sputtered bimetallic Cu-Co catalysts. In situ synchrotron X-ray absorption spectroscopy was employed to demonstrate dynamic stability of metallic Cu and Co phases under operating conditions. Besides, the authors further proposed synchrotron-based infrared spectroscopy to investigate the specific reaction mechanism. DFT calculations provide further evidence for the potential determining step. Finally, they achieve a high urea yield rate of 61 ± 6 mmol h‒1 gcat‒1 at –1.2 V vs. RHE. The work is well-structured, given the growing interest in sustainable nitrogen chemistry, and offers meaningful contributions to the understanding of C-N coupling mechanisms. However, before it could be accepted, some questions still need to answer.\n1. The tandem mechanism is attributed to atomic-level mixing via co-sputtering, but it is unclear whether such synergy is superior to a physical Cu + Co composite. Additional sample synthesized from physically mixed Cu/Co electrodes or sequential Cu-on-Co deposition should be discussed.\n2. The Co 2p XPS spectra appear to exhibit a broad peak that overlaps with a smaller underlying feature. The peaks should be re-fitted with appropriate constraints to ensure accurate deconvolution.\n3. The authors emphasize that neutral pH favors urea formation, yet the mechanistic rationale is only superficially discussed, which should be strengthened.\n4. It’s recommended to normalize current and yield rate to electrochemically active surface area (ECSA) to isolate intrinsic catalytic activity and decouple surface area effects.\n5. The authors seem to construct a Cu(111)/Co(111) plane model for DFT calculations. I am wondering whether the authors consider the defects or amorphous regions formed during sputtering, which may influence adsorption energetics.\n6. While stability tests are provided, the post-reaction TEM/XRD characterization is missing."}, "author_response": "Response to the first Reviewer’s comments\nComment: The authors developed a CuCo bimetallic catalyst for electrochemical synthesis of urea\nvia the co-reduction of CO and nitrite (NO ⁻). The catalyst was comprehensively characterized using\n2 2\na series of microscopic and spectroscopic approaches. In particular, in situ X-ray absorption\nspectroscopy (XAS), Raman, and infrared spectroscopy were employed to probe the dynamic chemical\nstate and adsorbed species under reaction conditions. These experimental results, combined with\ntheoretical calculations, provide insights into the reaction mechanism. Several control experiments\nfurther confirm that a Cu:Co ratio of 1:1 delivers the best performance in terms of urea Faradaic\nefficiency and yield. The authors also explored the pH effect, highlighting the role of proton balance\nin modulating surface *H coverage and enhancing proton-coupled electron transfer. Overall, this is a\nsystematic study for the newly developed catalyst material for the complicated electrochemical co-\nreduction of CO and NO ⁻. However, several key aspects regarding the nature of the active sites and\n2 2\nthe interaction between Cu and Co remain insufficiently explored. Moreover, some data were not\nproperly explained and clarified. Specific comments are listed below:\nReply: We sincerely thank the Reviewer for the careful evaluation of our manuscript and the\nconstructive comments provided. We have carefully revised the manuscript to address these concerns\nas outlined below:\n1. Introduction: The length of the introduction has been reduced to improve focus and\nreadability.\n2. Pure Cu performance: A detailed discussion on the poor performance of pure Cu toward\nelectrochemical C–N coupling has been added. Supporting evidence from in situ synchrotron\nFourier transform infrared (SR-FTIR) measurements has also been included.\n3. Effect of metal ratios: The interpretation of experimental results has been expanded by\ncorrelating Cu/Co ratios with electronic structure changes, and by discussing their influence on\nin situ SR-FTIR spectra and catalytic performance.\n4. XPS analysis: The fitting of Cu 2p and Cu XPS spectra has been revised and improved.\nLMM\n5. Cu–Co tandem redox: A comparative explanation of the redox behavior of Cu–Co relative to\npure Cu and pure Co has been incorporated, highlighting their synergistic role in enhancing\nelectrochemical C–N coupling.\n6. Presence of Co O in the catalyst: Extended Raman experiments were performed, confirming\n3 4\nthat increasing laser power contributes to the transformation of Co O from Co(OH) .\n3 4 2\n7. LCF of pure Cu: Linear combination fitting (LCF) has now been included for pure Cu to\nstrengthen the XAS analysis.\nWe trust that these revisions address the Reviewer’s concerns and significantly strengthen the\nmechanistic depth of the manuscript.\ncore message of the manuscript. Some contents are related to noble metal catalysts, which are\nirrelevant to this study as the manuscript focuses on non-noble metal catalysts.\nReply: We thank the reviewer for the valuable feedback regarding the length and focus of the\nIntroduction Part. In the revised manuscript, we have substantially condensed the Introduction to\nhighlight the core message of the study—namely, the rationale for using Cu–Co tandem catalysts for\nelectrochemical C–N coupling to synthesize urea, and the importance of in situ spectroelectrochemical\nanalysis for mechanistic understanding.\nIn addition, we have removed background content unrelated to the scope of this work, including\ndiscussions of noble metal catalysts. This streamlining ensures that the Introduction is more concise,\nfocused on non-noble metal systems, and directly aligned with the objectives and findings of the\npresent study.\nTherefore, we have made corresponding adjustments to the manuscript text in the main manuscript\n(Page 3-4, Lines 2-53) presented as follows:\n“Urea (CO(NH ) ) is the most utilized fertilizer globally with market size around $129 billion in 2023,\n2 2\nwhich it is projected to grow to around $161 billion by 2032.1 Currently, the primary method for\ncommercial urea production is the Bosch-Meiser process, which is involves the reaction of ammonia\n(NH ) and CO under elevated temperature (150 – 200 ⁰C) and pressure (100 – 200 bar).2 NH is\n3 2 3\ntypically derived from the Haber-Bosch process, which relies on hydrogen sourced from fossil fuels,\nmaking upstream hydrogen production the main contributor to energy consumption and carbon\nemissions.2 As a result, the overall urea production remains carbon-intensive, generating\napproximately 910 kg of CO for every ton of urea produced.3,4\n2 x 2 3\nsustainable alternative. These NO ⁻ species can be sourced from exhaust flue gas, chemical production\nx\nby-product, wastewater streams, or synthesized via cold plasma oxidation of N .5-10 Such approaches\n2\ncould enable decentralized and potentially carbon-neutral urea production. However, urea\nelectrosynthesis from CO and NO ⁻ is a complex multistep process involving multiple electron transfer\n2 x\n(Table S1-S2) and competing with mechanistically simpler reduction processes, viz. the hydrogen\nevolution reaction (HER), CO reduction reactions (CO RR), and NO ⁻/NO ⁻ reduction reactions\n2 2 2 3\n(NO ⁻RR/NO ⁻RR or NO ⁻RR).2,11 Notable challenges include limited CO adsorption capacity,12\n2 3 x 2\nsluggish dissociation of the C=O bond,12,13 and difficulty in balancing adsorption and activation of\nboth carbon- and nitrogen-derived species.2,14\nAt the core of these challenges lies the C–N coupling step, whose mechanism remains poorly resolved\ndue to interplay of multiple surface-bound species and dynamic changes in catalyst structure under\nreaction conditions. Proposed C–N coupling pathways involve intermediates such as *CO, *NO,\n*NH₂, and *NH, with key steps often limited by proton-coupled electron transfer (PCET) barriers.11,15-\n2\noperation—requires in situ spectroelectrochemical characterization techniques, such as Raman,\ninfrared, and X-ray absorption spectroscopy (XAS), which can directly capture transient intermediates\nand structural changes.12,13 Achieving the most energetically favorable pathway requires precise\ntuning of intermediate adsorption energies, a task that may be inherently challenging for a\nmonometallic catalyst. This has led to growing interest in multimetallic designs, where distinct active\nsites can preferentially stabilize different intermediates and facilitate sequential steps. Indeed, such\nstrategy has been successfully applied in the oxygen evolution reaction (OER)18,19 and CO RR20,21\n2\nfields, offering a blueprint for tackling the complexity of urea electrosynthesis.\nCopper has been widely studied due to its ability to facilitate CO and NO ⁻ activation, moderate\n2 x\nhydrogen binding, and tendency to form key intermediates such as *CO and *NH .22-26 However, Cu\n2\nsuffers from limited urea selectivity and deactivation caused by strong adsorption of nitrogen\nintermediates.27-31 Cobalt, on the other hand, offers high conductivity, favorable *CO binding, and the\nability to stabilize *H intermediates, thereby promoting downstream protonation or hydrogenation\nsteps.30,32-35 We hypothesize that combining Cu and Co could couple their complementary strengths—\nenhancing NO ⁻ reduction to reactive nitrogen species on Cu, while stabilizing *CO and *H on Co—\nx\nto accelerate C–N bond formation.30,36,37 Supporting this concept, recent work on amorphous Cu–Co\nboride (a-Cu CoB ) demonstrated alleviation of strong N-based and *CO intermediates binding on\n0.1 x\nthe catalyst surface and achieved a faradaic efficiency (FE) of ca 28% with a urea yield rate of ca 310\nμg h-1 mg -1 (5.2 mmol h-1 g -1) from CO and NO ⁻.38\ncat cat 2 3\nBuilding upon the hypotheses discussed above, we explored tandem Cu-Co catalysts prepared via co-\nsputtering of the two metals to facilitate C‒N coupling through the provision of different sites that\nfacilitate different key steps on this complicated process. We expected that the atomic-level mixing\nachieved through co-sputtering of Cu and Co can enhance PCET and facilitate the formation of key\nintermediates, thereby optimizing the C‒N coupling process.39-41 We first performed systematic\nelectrochemical studies under varying electrolyte conditions—including pH and nitrogen-source\nspecies—to evaluate urea formation and identify optimal conditions. Guided by these results, we\nemployed in situ spectroelectrochemical techniques—in situ Raman, synchrotron radiation Fourier-\ntransform infrared spectroscopy (SR-FTIR), and XAS—to monitor the formation and evolution of\nreaction intermediates directly on the catalyst surface under operating conditions. Finally, density\nfunctional theory (DFT) calculations were performed under the optimized reaction conditions to\nprovide atomic-scale insights into the roles of Cu and Co in facilitating key reaction steps and\nstabilizing crucial intermediates such as *CO, *NH , and *COOH. This integrated experimental–\n2\ntheoretical approach aims to unravel the C–N coupling mechanism and guide the rational design of\nefficient catalysts for sustainable urea electrosynthesis.”\nrecent studies showing urea production from CO and nitrate/nitrite using Cu or Cu O (Nat. Commun.\n2 2\n2025, 16, 3271; Applied Catalysis B: Environ. 2024, 340, 123265). The authors should provide more\ndetailed discussion or justification regarding this issue.\nReply: We appreciate the reviewer drawing attention to recent reports of urea formation on Cu/Cu O.\nmethod under our operating conditions, and we would like to reiterate to the Reviewer that this does\nnot contradict studies that employ oxide-derived Cu (OD-Cu) or Cu O under different regimes.16,42-45\n2\nreconstructs into oxide-derived Cu (OD-Cu)—metallic Cu enriched with residual Cu⁺ sites, subsurface\noxygen, defect structures, and increased surface roughness.42-44 Such features enhance *CO and\nnitrogen-intermediate co-adsorption, thereby promoting C–N bond formation.42 In contrast, the pure\nCu catalyst in this work was prepared by sputtering, which generates thin, relatively smooth films with\nfewer defects that weaken its ability to stabilize intermediates.45 Enhancement of C‒N coupling in\nsputtered Cu systems is therefore typically achieved by introducing a co-catalyst to create binary\nalloys, which tune the adsorption energetics and selectivity.16\nLinear combination fitting analysis of in situ XAS data for the pure Cu catalyst indeed indicated the\npresence of minor Cu O and Cu(OH) phases alongside metallic Cu. However, because linear\n2 2\ncombination fitting reflects bulk-averaged phase fractions, it cannot resolve whether these oxides are\nsurface-localized, catalytically accessible, or stable under reaction conditions.46,47 In parallel, our in\nsitu synchrotron Fourier transform infrared (SR-FTIR) measurements further reveal that the\nelectroreduction of CO and NO ⁻ on pure Cu is dominated by CO RR, showing intense *CO/*COOH\n2 x 2\nbands but only weak *CO (carbonyl) and *CN signals, characteristic of urea (Fig. R1, adapted as Fig.\nS42 in the revised Supporting Information). Together, these findings suggest that while trace Cu O\n2\nand Cu(OH) are present, their coverage and lifetime are insufficient to stabilize polar intermediates\n2\n(*NCO, *NH₂CO) necessary for effective C–N coupling. Consequently, the catalyst behaves\nessentially as clean metallic Cu, consistent with the absence of measurable urea formation in 1H-NMR\nor GC-MS analysis.\nTo clarify this point, we have added control experiments directly comparing in situ electroreduction\non Cu Co , pure Cu, and pure Co (Fig. R1) and expanded the discussion to explain the origin of this\n1 1\ndiscrepancy with literature results in detail and have added control experiments for clarity in main\nmanuscript (Page 10, Lines 208-218) and (Page 11 and Lines 232-236).\n“We also investigated electroreduction of CO and NO – at −1.2 V vs. RHE using monometallic Cu\n2 2\nand Co catalysts (Fig. S30 & Fig. S33) as control references. The sputtered Cu catalyst produced no\ndetectable C–N products, instead yielding formate, ethanol, and propanol (Fig. S29, S47-S48). This\noutcome contrasts with many reports describing strong C–N coupling on Cu and underscores a more\nnuanced picture.15,48 In the literature, high urea selectivity is typically associated with oxide-derived\n*NCO and *NH CO.42,44 By comparison, sputtered Cu typically generates thin, relatively smooth films\n2\nwith fewer defect sites, thereby limiting the stabilization of key intermediates required for C‒N\ncoupling.45 Consequently, metallic Cu favors CO adsorption and *CO/*COOH formation,49\n2\nconsistent with the results observed in the in situ SR-FTIR spectra (Fig. S42). For pure Co, the main\nproducts were ammonium and H (Fig. S33), with no evidence of C–N products, also corroborated by\n2\nin situ SR-FTIR (Fig. S42).\nThese controls underscore that neither Cu nor Co alone is sufficient for effective C–N coupling.\nInstead, the tandem interface of Cu–Co is essential: Cu sites efficiently adsorb and activate CO to\n2\n*CO, while Co sites accelerate the stepwise reduction of NO ⁻ to *NH /*NH . The spatial co-\n2 2 3\ngeneration and proximity of *CO and *NH at the Cu–Co perimeter greatly enhances the likelihood\n2\nof C–N bond formation, leading to markedly higher urea activity and selectivity than either metal\nalone.”\n2 2\nreduction using (A) Cu Co (B) Cu, and (C) Co catalysts. The electroreduction of CO and NO ⁻ was carried out in the\n1 1, 2 2\nCO -saturated 0.2 M KHCO and 0.02 M KNO solution. For the Cu catalyst, CO reduction reaction predominates, with\n2 3 2 2\ncarbonyl (*CO) and *CN bands associated with urea observed at approximately 1590 cm⁻1 and 1455 ⁻1, respectively, though\ntheir intensities are weaker than those of *COOH and *CO species related to formate (~1720 cm⁻1) and gaseous CO (~2065\ncm⁻1). In contrast, no carbonyl or *CN features were detected for the Co catalyst.\nproduction. However, the detailed interaction between Cu and Co was not well demonstrated in the\nmanuscript. In addition, the linear combination analysis of XAS data shows the catalyst is not fully\nreduced under negative applied potential, with the co-existence of oxidized Cu or Co species. Their\namount is also varied at different reaction conditions. Therefore, the oxidized species might play an\nimportant role for catalysis, which should also be further explored.\nReply: We thank the reviewer for this insightful comment. To better understand the role of each Cu\nand Co, we conducted additional in situ SR-FTIR measurements on pure Cu and pure Co catalyst as\ncontrols (Fig. R1, depicted as Fig. S42 in the revised Supporting Information). The results shows that\npure Cu exhibits higher activity towards CO RR, as evidenced by the formation of *CO bands\n2\nassociated with CO , *COOH associated with formate, and *CO stretching related to methanol\n(g)\nformation.50-52 Carbonyl (*CO) and *CN bands related to urea were only very weakly detected,\nindicating that urea concentration was below detection limit of 1H-NMR or GC-MS. In contrast, in situ\nspectra for pure Co, revealed no *CN features. Only weak *CO/*COOH bands were observed, while\nstrong signals corresponding to *NH were detected, suggesting dominant activity for NO ⁻RR (Fig.\n3 2\nR1). These controls indicate that while both Cu and Co individually show activity for CO RR and\nminor contributions from Cu O/Cu(OH) and Co O /Co(OH) . Such oxides/hydroxides have indeed\n2 2 3 4 2\nbeen reported to stabilize polar intermediates (e.g.,*NCO)53-55 and to promote NO ⁻ activation and\n2\nhydrogenation, repectively.30,56 However, when compared with the in situ SR-FTIR results for pure\nCu, only weak carbonyl and *CN bands were observed. It is clear that these minority oxide/hydroxide\nphases, although present, are insufficient in coverage and/or lifetime to stabilize the intermediates\nnecessary for effective C–N coupling. Moreover, our CV measurements (Fig. R2, depicted as Fig. S24\nin the revised Supporting Information) also suggest the reduction of Cu and Co into their metallic\nstates at applied potential more negative than −1 V vs. RHE. Since linear combination fitting XAS\nprovides only bulk-averaged phase fractions and cannot resolve whether these minor phases are\nsurface-localized, catalytically accessible, or persistent under reaction conditions,57 we regard these\nvalues as indicative rather than definitive.\nCollectively, the evidence indicates that the predominant active states under cathodic bias are metallic\nCu and Co. Thus, urea formation is governed not by the minor oxidized species, but by the tandem\nCu–Co interface, where Cu facilitates *CO/*COOH generation and Co promotes *NH formation in\n2\nclose proximity—greatly increasing the likelihood of C–N bond coupling. This cooperative behavior\nlikely arises from electronic and structural interactions at the Cu–Co interface. Supporting this,\nelectrochemical impedance spectroscopy (EIS) reveals that Cu Co exhibits a lower charge transfer\n1 1\nresistance than pure Cu or Co (Fig. R3, depicted as Fig. S27 in the revised Supporting Information),\nunderscoring the enhanced electronic synergy that facilitates CO and NO ⁻ and underpins the C–N\n2 2\ncoupling performance of the bimetallic catalyst.\nsimple drop-casting (Fig. R4, depicted as Fig. S33 in the revised Supporting Information). No urea\nwas detected under identical conditions, underscoring that atomic-level mixing achieved through co-\nsputtering is essential for enabling C–N coupling.\nTo clarify this point, we have expanded the discussion in the main manuscript in Page 10, Lines 208-\n218; Page 11-12, Lines 232-263; and Page 16-17, Lines 385-398, as follows:\nPage 10, Lines 208-218; Page 11-12, Lines 232-263\n“We also investigated electroreduction of CO and NO – at −1.2 V vs. RHE using monometallic Cu\n2 2\nand Co catalysts (Fig. S30 & Fig. S33) as control references. The sputtered Cu catalyst produced no\ndetectable C–N products, instead yielding formate, ethanol, and propanol (Fig. S29, S47-S48). This\noutcome contrasts with many reports describing strong C–N coupling on Cu and underscores a more\nnuanced picture.15,48 In the literature, high urea selectivity is typically associated with oxide-derived\nCu, where mixed Cu(0)/Cu(I) and defect-rich surfaces stabilize polar intermediates such as *CO,\n*NCO and *NH CO.42,44 By comparison, sputtered Cu typically generates thin, relatively smooth films\n2\nwith fewer defect sites, thereby limiting the stabilization of key intermediates required for C‒N\ncoupling.45 Consequently, metallic Cu favors CO adsorption and *CO/*COOH formation,49\n2\nconsistent with the results observed in the in situ SR-FTIR spectra (Fig. S42). For pure Co, the main\nproducts were ammonium and H (Fig. S33), with no evidence of C–N products, also corroborated by\n2\nin situ SR-FTIR (Fig. S42).\nThese controls underscore that neither Cu nor Co alone is sufficient for effective C–N coupling.\nInstead, the tandem interface of Cu–Co is essential: Cu sites efficiently adsorb and activate CO to\n2\n*CO, while Co sites accelerate the stepwise reduction of NO ⁻ to *NH /*NH . The spatial co-\n2 2 3\ngeneration and proximity of *CO and *NH at the Cu–Co perimeter greatly enhances the likelihood\n2\nof C–N bond formation, leading to markedly higher urea activity and selectivity than either metal\nalone.\nElectrochemical CV measurements further highlight this synergy, as the Cu–Co bimetallic catalyst\nexhibits intermediate quasi-reversible behavior between pure Cu and Co (Fig. S24). This suggests\nthat Cu suppresses Co dissolution by stabilizing its redox transitions, while Co mitigates Cu over-\noxidation.58-60 Such mutual stabilization likely arises from electronic and structural interactions at the\nCu–Co interface, further reinforcing the unique cooperative role of the tandem system. Consistently,\nelectrochemical impedance spectroscopy (EIS) reveals that Cu Co has a lower charge transfer\n1 1\nresistance compared to pure Cu and Co (Fig. S27), underscoring its enhanced electronic synergy that\nfacilitates CO and NO ⁻ and underpins the C–N coupling performance of the bimetallic catalyst.\n2 2\nImportantly, this effect cannot be replicated by simple physical mixing: a drop-cast Cu/Co composite\nproduced only ammonium, formate, H , CO, and C alcohols (Fig. S33), with no detectable urea.\n2 2+\nThis finding demonstrates that atomic-level proximity achieved by co-sputtering is critical for\ncooperative interactions and the relay-type mechanism required for urea synthesis.\nin the Cu /Co matrix on the FE and yield rate of urea. As compared to Cu Co , the Cu Co\n(n) (1- n) 1 1 2 1\ncomposition produces lower j at potentials more negative than −1.0 V vs. RHE (Fig. 4D). Methylamine\nis detected as the main C‒N product, albeit at a low FE of less than 1%, within the potential range of\n−1.0 to −1.2 V vs. RHE (Fig. 4E). The increased amount of Cu in the catalyst favored the C‒C\ncoupling, producing C products such as ethanol and n-propanol. At −1.2 V vs. RHE, the FEs for\n2+\nethanol and n-propanol achieve 22 ± 4% and 2.1 ± 0.4 %, respectively. At −1.3 V vs. RHE, the FE of\nboth products improved to 25 ± 2% and 2.4 ± 0.7%, accordingly. Conversely, the use of the Co-rich\nCu Co catalyst enhances j for the electroreduction of CO and NO ⁻ at potentials more negative than\n1 2 2 2\n−1.0 V vs. RHE as compared to Cu Co (Fig. 4D). However, neither urea nor methylamine were\n1 1\nproduced. The predominant products from the Cu Co catalyst are ammonium, formate, and H . These\n1 2 2\nresults highlight the importance of a specific Cu:Co ratio in balancing the adsorption and reduction\nof reactants, thus promoting urea generation without overshadowing the process with excessive side\nreactions (HER, CO RR, and NO ⁻RR). Moreover, all Cu Co catalysts demonstrated good\n2 2 (n) (1-n)\nstability over the course of 12 h chronoamperometric experiments for the co-reduction of CO and\n2\nnitrite, as confirmed by ICP-MS analysis (Fig. S35), showing less than 10 wt% loss of Cu and Co.\nThis loss is considered low compared to the literature.29,61”\nPage 16-17, Lines 385-398\n“Together, these results establish a clear structure-activity correlation. Minor copper\noxides/hydroxides detected by LCF EXAFS/XANES may, in principle, stabilize polar intermediates\nsuch as *NCO.53-55 Similarly, cobalt oxides/hydroxides have been reported to enhance NO ⁻ activation\n2\nand hydrogenation.30,56 However, these species are present only in trace amounts, and their limited\nsurface coverage or transient lifetime appears insufficient under electroreduction conditions, as\ndepicted in CV measurements (Fig. S23). This is consistent with the in situ SR-FTIR spectra of pure\nCu, which show only weak *CN features (Fig. S42). It should also be noted that LCF provides bulk-\naveraged phase fractions and cannot resolve whether these minority species are surface-localized,\ncatalytically accessible, or stable under bias; hence, their role is considered indicative rather than\ndefinitive for the active surface. Overall, the evidence indicates that Cu alone enhances CO\n2\nadsorption and stabilizes *CO/*COOH intermediates but lacks sufficient *NH coverage; Co alone\n2\nreduces NO ⁻ to *NH /*NH but provides inadequate *CO (Fig. S42). In contrast, at the Cu–Co\n2 2 3\ninterface, CO -derived *CO/*COOH from Cu and NO ⁻-derived *NH from Co are generated in close\n2 2 2\nproximity, which greatly enhances the probability of C–N bond formation. Guided by these\nmechanistic insights, we then performed DFT modeling of the reaction pathways.”\n1 1 2 2\nelectroreduction of CO and NO ⁻. CV curves for (A-C) HER, (D-F) CO RR, (G-I) NO ⁻RR, and (J-L) electroreduction\n2 2 2 2\nof CO and NO ⁻. CV measurements show that pure Cu exhibits reversible redox behavior (two oxidation and two reduction\n2 2\npeaks), whereas pure Co displays irreversible features (only one oxidation peak), attributed to Co2+/Co3+ dissolution into\nthe electrolyte.62,63 In contrast, the Cu–Co tandem stabilizes Co, suppresses dissolution, and modifies its electronic\nstructure, yielding a quasi-reversible system with enhanced stability and reusability in electrochemical reactions.58-60\nAll experiments were conducted at scan rate of 20 mV s⁻1 and scan number = 5. The CV measurements for HER, CO RR,\n2 2 2 3 2\n0.2 M KHCO solution, Ar-saturated 0.2 M KHCO + 0.02 M KNO solution, and CO -saturated 0.2 M KHCO + 0.02 M\n3 3 2 2 3\nKNO solution, respectively.\n2\n(n) (1-n)\nand (B) the circuit model described the electrochemical system used to fit the raw data. Panel (A) shows the Nyquist plots\nsemicircles, representing response to varying frequency at a fixed potential. Panel (B) provides the equivalent circuit to fit\nthe Nyquist plots obtained, where R is the solution electrical or electrolyte resistance that can also represents distance\ns\nbetween working electrode and counter electrode, R is the charge transfer resistance that represents interfacial resistance\nct\nbetween electrode and electrolyte, and constant angle element (CPE) is an empirical constant that mimics non-ideal\ncapacitance. R is constant for all catalyst due to the distance between working and counter electrode being kept constant.\ns\nAs can be seen, when ratio between Cu and Co is equal, the R achieve the lowest value, indicating lower contact resistance\nct\nand work function, thereby shorten diffusion pathway for electrolyte.64,65\n2 2\n(A) Cu and Co catalysts (prepared via sputtering), a physically mixed Cu/Co catalyst (prepared by drop-casting), and (B)\ndifferent ratio of co-sputtered Cu Co catalysts. All experiments were carried out in CO -sat. 0.2 M KHCO + 0.02 M\n(n) (1-n) 2 3\nKNO solution at -1.2 V vs. RHE. The Cu catalyst exhibits activity towards CO RR to generate C products (formate,\n2 2 1\nmethane, and methanol) and C products (acetate, ethanol, and propanol), whilst Co catalyst has activity towards\n2+\nconversion of NO ⁻-to-NH + and H with trace of formate, and CH . The urea is only detected when combining Cu and Co\n2 4 2 4\ncatalyst with equal ratios. Increasing Cu ratio on Cu Co catalysts enhanced the activity towards C products.\n(n) (1-n) 2+\nFor the drop-cast Cu/Co control, the catalyst ink was prepared by mixing 25 μmol CuO (Sigma Aldrich, 99.99%), 25 μmol\nCoO (Sigma Aldrich, 99.5%), 6 mg carbon powder, 500 μL ethanol, 500 μL water, and 30 μL Nafion solution, followed\nby deposition onto the electrode substrate.\nNote: H = hydrogen, CO = carbon monoxide, CH = methane, NH ⁺ = ammonium, HCOO⁻ = formate, CH COO⁻ = acetate,\n2 4 4 3\nCH OH = methanol, C H OH = ethanol, C H OH = propanol, CO(NH ) = urea, and CH NH = methylamine.\n3 2 5 3 7 2 2 3 2\nconclusion of existing of Cu(0), Cu(II), but not Cu(I), is questionable. The peak deconvolution of\nCuLMM Auger spectra in Fig. S10e is also not accurate. Please check the relevant literatures (e.g.\nSurf. Interface Anal. 2017, 49, 1325) for more accurate XPS data fitting.\nReply: We thank the reviewer for highlighting the limitations of Cu 2p analysis. We agree that Cu 2p\nalone cannot unambiguously separate Cu(0) from Cu(I). Accordingly, we have now re-analyzed the\nspectra following Biesinger’s protocol,66 combining Cu 2p with Cu Auger analysis and the\nLMM\nmodified Auger parameter (α′ = BE[Cu 2p ] + KE[Cu ]) to determine the presence of\n3/2 LMM\nCu(0)/Cu(I)/Cu(II). Moreover, we also quantified the composition based on the fitting of Cu\nLMM\nregions.\nFollowing the Reviewer’s feedback, we deconvoluted the Cu(0), Cu(I), and Cu(II) components for all\nCu Co catalysts (Fig. R5-R7; adapted as Fig. 2, Fig. S11, and Fig. S13 in the revised Supporting\n(n) (1-n)\nInformation). Cu(II) is evident from the pronounced shake-up satellite structure in Cu 2p region, and\nthe satellite pattern is consistent with Cu(OH) -like species reported by Biesinger.66 Quantification\n2\nfrom the Cu spectra shows that the as-prepared Cu Co catalyst contains a mixture of oxidation\nLMM 1 1\nstates: ~32% Cu(0), 30% Cu(I), and 38% Cu(II) (Fig. R5). We further compared the modified Auger\nparameter of our samples with reference value (Table R1, adapted as Table S4 in the revised\nSupporting Information), confirming mixed Cu(0)/Cu(I)/Cu(II) speciation in the catalyst.\nWe have incorporated this revised analysis in the main manuscript (Page 6, Lines 82-97):\n“X-ray photoelectron spectroscopy (XPS) was employed to elucidate the surface chemical composition\nof the as-prepared Cu Co catalysts (Figs. 2A-B & Figs. S9-S14), with monometallic sputtered\n(n) (1- n)\nCu and Co serving as references. The survey spectrum reveals distinct peaks corresponding to C 1s,\nO 1s, Cu 2p and Co 2p, with relative intensities of Cu 2p and Co 2p peaks following the bulk\ncomposition (Fig. S9). Fitting of the high-resolution Cu 2p and Co 2p spectra of as-synthesized\ncatalysts using the parameters reported by Biesinger et al.67,68 and our previous reports,69,70 as a\nguide displays the presence of Cu(0), Cu(I), Cu(II), Co(0), and Co(II) species, indicating oxidation of\nsurface Cu and Co metal to Cu(I), Cu(II) and Co(II) due to air exposure (Fig. 2A-B, Figs. S11-S12,\n& Table S4).71,72 High-resolution Cu Auger spectra further confirms the presence of Cu(0), Cu(I),\nLMM\nand Cu(II) species (Fig. S13). Analysis of the high-resolution O 1s spectra of Cu Co catalysts\n(n) (1- n)\nidentifies lattice (529.3 eV) and non-lattice oxygen (~531.5 and 532.5 eV) (Fig. S14).73 Post-reaction\nanalysis of Cu Co catalyst reveals changes in surface structure after 12 h electrochemical testing\n(n) (1- n)\n(vide infra) (Figs. 2A-B & Figs. S9). In addition to Cu 2p, Co 2p, C 1s, and O 1s peaks, survey spectra\ndetect K 2p and K 2s peaks, indicating adsorption of K+ from the electrolyte on the catalyst surface.\nAnalysis of Co 2p, Cu 2p, and Cu Auger spectra indicate further oxidation of Cu(0) and Co(0) to\nLMM\nCu(I), Cu(II) and Co(II) induced by the electrocatalytic testing (Figs. S11-S13).74 Consequently, the\noxidation leads to the increased peak intensity of lattice oxygen, as seen in high-resolution O 1s spectra\n(Fig. S14).”\n1 1\ndeposited catalyst. (C-D) The X-ray absorption near edge structure (XANES) at Cu K-edge (C) and Co K-edge (D). (E-F)\nThe k3-weighted EXAFS spectra at Cu K-edge (E) and Co K-edge (F). (G-H) The normalized FT-EXAFS of Cu Co at Cu\n1 1\nK-edge (G) and Co K-edge (H) of Cu Co catalyst. The catalysts were tested for 12 h under condition as in Fig.3.\n1 1\n(n) (1-n) 1 2\n(C-D) Cu Co , and (E-F) Cu Co catalyst. Cu 2p signals at 932.6, 932.2, and 934.6 eV in the pre- and post-reaction catalysts\n1 1 2 1 3/2\nrepresent Cu(0), Cu(I), and Cu(II), respectively,66,67 with the Cu2+ satellite consistent with Cu(OH) .66 The pre-reaction\n2\ncatalyst refers to the as-prepared catalyst, whereas the post-reaction catalyst refers to the catalyst after a 12 h reaction in\nthe CO -saturated 0.2 M KHCO + 0.02 M KNO .\n2 3 2\nLMM (n) (1-n) LMM 1 2\n(A) pre-reaction and (B) post-reaction. High-resolution XPS Cu Auger spectra for Cu Co : (C) pre-reaction and (D)\nLMM 1 1\npost-reaction. High-resolution XPS Cu Auger spectra for Cu Co : (E) pre-reaction and (f) post-reaction. The curve-\nLMM 2 1\nfitting of the Cu Auger spectra elucidate Cu(0), Cu(I), and Cu(II) species.66,67 The Cu (0) signals weaken after the\nLMM\nreaction, indicating oxidation of Cu(0) to Cu(I)/Cu(II). The pre-reaction catalyst refers to the as-prepared catalyst, whereas\nthe post-reaction catalyst refers to the catalyst after a 12 h reaction in the CO -saturated 0.2 M KHCO + 0.02 M KNO .\n2 3 2\n3/2 LMM (n) (1-n)\ncatalyst. The calculation of Auger parameter was based on Biesinger’s method (α′ = BE[Cu 2p ] + KE[Cu ]).66\n3/2 LMM\nNo Sample Refer. Cu 2p Cu Auger Ref.\n3/2 LMM\nComp. peak max peak max parameter Auger\nEb (eV) Ek (eV) α′ (eV) parameter\nα′ (eV)\nCu(0) 932.66 918.49 1851.15 1851.24\n1 1\nCu(II) 934.70 916.38 1851.08 1850.92\nCu(0) 932.73 918.58 1851.31 1851.24\n1 2\nCu(II) 934.77 916.41 1851.18 1850.92\nCu(0) 932.63 918.10 1850.73 1851.24\n2 1\nCu(II) 934.66 916.36 1851.02 1850.92\nHowever, the detailed information of the reaction condition was not clearly provided. The phenomena\nof further oxidation of the catalysts compared to the pre-catalysts after electrochemical reduction\nconditions are strange. It is highly probable that the catalyst materials were extensively oxidized in air\nafter electrochemical tests. Given that the authors performed in situ experiments in the following\nsection, the post-reaction characterization seems not necessary and it may mislead the readers\nregarding the real active species.\nReply: We thank the reviewer for this valuable comment. Post-reaction characterization was\nperformed on catalysts after 12 h of electroreduction in CO -saturated 0.2 M KHCO + 0.02 M KNO .\n2 3 2\nAfter electrolysis, the samples were stored in a vacuum chamber to minimize oxidation. Nevertheless,\nwe acknowledge that some re-oxidation is unavoidable during sample handling, transport, and\nanalysis. Even so, EXAFS and its LCF analyses of the post-reaction catalysts indicate that metallic Cu\nand Co remain the dominant phases.\nAs also requested by Reviewer 3, we have presented post-reaction analyses to provide complementary\ninsights into catalyst stability and re-oxidation behavior. We fully agree that in situ techniques provide\nthe most reliable information on the active state of the catalyst under electrochemical conditions.\nAccordingly, we have emphasized the in situ results as the primary evidence in our mechanistic\ndiscussion. At the same time, we retained post-reaction characterizations (XAS and XPS) as\ncomplementary information to provide insights into catalyst stability, the extent of re-oxidation after\nelectrolysis, and the reversibility of structural changes.\nTo avoid potential misunderstanding, we have clarified in the revised manuscript that the post-reaction\ncharacterizations are presented in the Supporting Information solely for completeness and to illustrate\nre-oxidation behavior, but not as evidence of the active species under bias.\nIn addition, we have now included further details of the post-reaction characterization procedure in the\nmain manuscript (Page 23, Lines 566-568):\n“For post-reaction characterization, the catalysts were subjected to 12 h of electroreduction in CO -\n2\nsaturated 0.2 M KHCO + 0.02 M KNO solution. After electrolysis, the samples were collected, rinsed\n3 2\nwith deionized water, and stored in a vacuum chamber until further analysis.”\nfor clarity.\nReply: We thank the reviewer for this helpful suggestion. In the revised manuscript, we have carefully\ncondensed the section from Lines 97–112 to remove repetitive statements while retaining the essential\nresults. Specifically, we streamlined the description of the XANES, EXAFS, FT-EXAFS, and LCF\nanalyses into a single coherent narrative, avoiding overlap in reporting the metallic states of Cu and\nCo. The revised text now highlights the key findings—that the as-prepared catalysts are dominated by\nmetallic Cu and Co with minor oxide/hydroxide contributions, and that post-reaction samples remain\nprimarily metallic with small fractions of oxidized species—presented with improved clarity in the\nmain manuscript (Page 6, Lines 98-109).\n“XAS measurements of Cu Co catalysts were recorded at Cu and Co K-edges. The X-ray\n(n) (1-n)\nabsorption near edge structure (XANES) and extended X-ray absorption fine structure (EXAFS)\nspectra (Fig. 2C-H & Fig. S15-S18) confirm that the as-prepared catalysts are dominated by metallic\nCu(0) and Co(0), consistent with the XRD and HRTEM results. Fourier transformation (FT-EXAFS)\nand k3-weighted EXAFS spectra further support this assignment (Fig. S16). Linear combination fitting\n(LCF) of both XANES (Figs. S19) and EXAFS (Figs. S20) indicate that the materials being composites\nprimarily consisting of metallic Cu and Co, without evidence of intermetallic phases. In addition to\nthe dominant metallic phases, minor Cu O/Cu(OH) and Co O /Co(OH) species are also present as\n2 2 3 4 2\ndetectable oxide/hydroxide phase. After electroreduction, XANES spectra (Fig. S15, S17–S18) show\nslight oxidation, evidenced by an increase in edge intensity, consistent with XPS results. Nevertheless,\nFT-EXAFS at both Cu and Co K-edges confirms that Cu–Cu and Co–Co metallic interactions remain\npredominant (Fig. S17–S19), with LCF analysis indicating that the post-reaction catalysts are still\nprimarily metallic, with fractions of oxidized Cu and Co species present (Fig. S19-S20).”\n3 4\nto laser-induced oxidation. However, this statement might be incorrect, as XAS is more bulk sensitive\nand may not detect near surface Co O , particularly in small quantities. The LCA data in Fig. S20 also\n3 4\nshows the presence of Co O for the other control CuCo samples, which may have higher amount of\n3 4\nCo O .\n3 4\nReply: We thank the Reviewer for the helpful clarification. We agree that bulk-sensitive XAS may\nmiss a very thin, near-surface Co O fraction. Our XAS linear-combination fitting places an upper\n3 4\nbound of <8% Co O in the as-prepared Cu Co (as shown in Fig. S19-S20 in the revised Supporting\n3 4 1 1\nInformation), indicating that any pre-existing Co O is minor and may fall below XAS bulk\n3 4\ndetectability.\nTo probe the Raman/XAS discrepancy, we performed power-dependent Raman controls on Cu Co ,\n1 1\npure Cu, and pure Co (Fig. R8, adapted as Fig. S21 in the revised Supporting Information). At 50%\nlaser power, pure Cu shows Cu O features and pure Co shows Co(OH) ; the Cu Co sample likewise\n2 2 1 1\ndisplays Cu O and Co(OH) . Upon increasing laser power above 50%, additional oxide bands appear:\n2 2\nCu evolves to CuO/Cu(OH) , Co evolves to Co O , and Cu Co exhibits mixed of\n2 3 4 1 1\nCuO/Co O /Co(OH) . This systematic, power-dependent growth of metal oxidation states is more\n3 4 2\nconsistent with a photo-oxidation contribution under Raman illumination.75-78\nComplementary XPS analysis does not resolve distinct Co O component in our samples (Fig. R5 &\n3 4\nR25, adapted as Fig. 2 & Fig. S12 in the revised manuscript and Supporting Information, respectively),\nand the Co 2p main peak and satellite structures are characteristic of Co(II), in agreement with the\nliterature.68 Taken together, Raman (with power controls), XPS, and linear combination fitting XAS\nindicate that any pre-existing Co O in our catalyst is at most a minor near-surface fraction; the Co O\n3 4 3 4\nbands observed by Raman can thus be rationalized by laser-induced oxidation, possibly superimposed\non a small surface oxide that lies below XAS detectability.\nBased on Reviewer’s suggestion, we revised the main manuscript as follow (Page 6-7, Lines 110–\n119):\n“Raman spectroscopy was also conducted to characterize the components of the Cu Co catalysts\n(n) (1- n)\n(Fig. S21). At 50% laser power, the as-prepared Cu Co catalyst exhibits peaks indicative of Co(OH)\n1 1 2\nat ~530 and 600 cm⁻1, Cu O at ~147 and 218 cm⁻1, and CuO at ~294 cm⁻1.77,79-82 These oxide/hydroxide\n2\nfeatures are attributed to the oxidation of metallic Cu and Co resulting from air exposure.83,84 Under\npower-dependent Raman acquisition, additional CuO and Co O modes emerge only at higher laser\n3 4\npower (>50%) that may indicate a laser-induced oxidation component.75,78 Meanwhile, LCF of EXAFS\nspectra (Fig. S20) place an upper bound of ≤8% Co O in Cu Co catalysts, implying that any\n3 4 (n) (1- n)\npre-existing Co O is minor near-surface fraction below bulk XAS detectability. Taken together,\n3 4\npower-dependent Raman and LCF EXAFS spectra support that the Co O bands observed by Raman\n3 4\narise primarily from laser-induced oxidation, possibly superimposed on a very thin surface Co O\n3 4\nlayer.”\n1 1\nvarying laser power. (D) Comparison of as-prepared Cu Co catalysts acquired at 100% laser power. For pure Cu, at\n(n) (1-n)\n50% laser power, Cu O bands are observed at ~147, 219, and 634 cm⁻1.30,80 Upon increasing laser power, additional copper\n2\noxides/hydroxide features emerge. A similar trend is observed for pure Co; at 50% laser power, Co(OH) bands are\n2\ndiscernible at ~467, 516, 602, and 666 cm⁻1.77,81,82,85 However, with higher power, the spectrum evolves to Co O ,\n3 4\nshowing modes at ~189, 469, 510, 610, and 675 cm⁻1.77,81,82,85 The emergences of oxides/hydroxides in higher laser power may\nreflect improved signal-to-noise and/or laser-induced oxidation during Raman acquisition.77,81,82,85\nFor the as-prepared Cu Co catalyst, spectra acquired at 50% laser power are dominated by Cu O and Co(OH) , whereas\n(n) (1-n) 2 2\nhigher laser power reveals CuO bands at ~276 and 320 cm⁻1.80 At 100% laser power, Cu Co exhibits a mixture of Co O\n1 1 3 4\nand Co(OH) features at ~189, 470, 512, 600, and 664 cm⁻1,77,81,82,85 together with CuO bands at ~276 and 320 cm⁻1.80\n2\n(Peak positions may vary by ±5–10 cm⁻1 due to local heating and instrumental parameters).\n2000–2100 cm⁻¹ can also correspond to *NO (see ACS Energy Lett. 2022, 7, 1187). This is because\nCO and NO are chemically similar. Isotope labeling is strongly recommended to accurately assign\nand distinguish these species.\nReply: We thank the Reviewer for this valuable comment and for pointing out the possibility that\npeaks between 2000–2100 cm⁻¹ may also correspond to *NO species, as reported in literature.86 We\nagree that *CO and *NO can have overlapping vibrational frequencies due to their chemical similarity.\nTo minimize the risk of peak misidentification, we have now performed isotope tracing during in situ\nsynchrotron Fourier-transform infrared (SR-FTIR) measurements using 15NO ⁻ as control experiment.\n3 2\nyields. Therefore, prior to isotope tracing, we compared the performance of 14NO ⁻ and 14NO ⁻ and\n2 3\nfound that both nitrogen sources yielded similar spectra (Fig. R9, adapted as Fig. S43 in the revised\nSupporting Information). On this basis, we proceeded with isotope tracing using 15NO ⁻.\nintermediates. However, the peak around 2065 cm⁻1 remained unchanged, indicating that it originates\nfrom carbon-containing intermediates, which we assign to *CO species related to CO gas formation.\nUnfortunately, we were unable to perform comparable analysis using in situ Raman spectroscopy.\nUnder our measurement conditions, no distinct peaks—including metal–surface-related signals—were\ndetected. We attribute this to differences in instrumentation and facility configurations relative to our\nprevious Raman experiments at the UNSW Mark Wainwright Analytical Centre Spectroscopy\nLaboratory, combined with the inherently weak Raman scattering of the catalyst under electrochemical\noperation. Consequently, reliable Raman data could not be obtained for presentation.\nBased on Reviewer’s suggestion, we also have incorporated the new founding in the main manuscript\n(Page 15, Lines 347-351):\n“Isotope labelling with 15NO ⁻ confirmed the nitrogen origin of the *CN band (Fig. S43). Spectra for\n3\n14NO ⁻ and 14NO ⁻ were nearly identical, apart from an additional *NO ⁻ feature at ~1340 cm⁻1.87\n2 3 3\nSubstitution with 15NO ⁻ shifted nitrogen-containing vibrational modes to lower wavenumbers, with\n3\nthe *C15N feature of urea detected at 1411 cm⁻1, while the carbonyl peak remained unchanged.88 This\nisotopic shift conclusively verifies that the *CN signal originates from NO ⁻-derived nitrogen.”\nx\n1 1\nCO and NO⁻ reduction using (A) KNO , (B) KNO , and (C) K15NO as nitrogen sources. Electroreduction was carried\n2 x 2 3 3\nout in the CO -saturated 0.2 M KHCO and 0.02 M KNO solution. K15NO was employed for isotope tracing, as KNO\n2 3 x 3 3\nand KNO yield comparable products under the tested conditions. Isotopic substitution with 15N shifts the corresponding\n2\nvibrational bands to lower wavenumbers, consistent with increased molecular mass.88\nas the activity of urea production might be over-estimated due to the influence of the contaminated\nspecies. Although in Fig. S32 the authors showed the 15N labelling NMR data, it would be better to\nquantify it to demonstrate that urea is indeed generated from nitrite.\nReply: We thank the Reviewer for emphasizing the importance of the 15N isotope labelling experiment\nin confirming the origin of urea. In response, we have now quantified the 15N-labelling NMR results\nshown in Fig.R10 (adapted as Fig. S32B-C in the revised Supporting Information), using the same\n1H-NMR protocol as 14N-urea. In 15N-labeled samples, the urea –NH₂ resonance appears as a 15N-\ncoupled doublet; we integrate the total area, which yields the same response factor as the 14N-urea\ncalibration within experimental error.\nBecause isotope tracing was performed with K15NO , we used K14NO as the direct comparison under\n3 3\nidentical conditions. With K14NO , the urea signal appears as a singlet at ~5.62 ppm and gives a yield\n3\nrate of 18 ± 2 mmol h⁻1g ⁻1. With K15NO , the urea ‒NH resonance appears as 15N-coupled doublet\ncat 3 2\nat ~5.58 ppm and ~5.71 ppm (consistent with reported literature).89 Urea quantification gives a yield\nrate around 16 ± 1 mmol h⁻1g ⁻1. The small difference between 14N and 15N runs lies within\ncat\nexperimental uncertainty and may also reflect minor isotope/relaxation effects or the exclusion of any\nunlabeled background nitrogen.89 Importantly, the isotopic shifts unambiguously confirm that the\ndetected urea derives from the NO ⁻ nitrogen source.\nx\nThe quantified data have been added to the revised Supporting Information (Fig.R10), and the\ncorresponding description has been incorporated into the main manuscript (Page 8-9, Lines 162-164):\n“Furthermore, 15N isotope-tracing experiments confirm that the produced urea originates from the\nNO ⁻ feed (Fig. S32 & Fig. S42).”\nx\nfragmentation pattern of the sample displays urea-specific fragment ions (m/z = 73, 99, 117, 132, 147, 157, 213, 231, and\n273). (B-C) The 1H-NMR spectrum of a sample after electroreduction of (B) CO + 14NO ⁻, (C) CO + 15NO ⁻, and (D)\n2 3 2 3\nCO + 14NO ⁻ (after water peak removal). (F) The 1H-NMR spectrum of ammonium carbamate (chemical shift = 7.4 ppm)\n2 2\nbecause of electroreduction of CO and NH OH. All experiments utilized 0.2 M KHCO as buffer solution and using\n2 2 3\napplied potential of -1.2 V vs. RHE for 12 h reaction.\nelectrolyte to interaction between the reactants to Cu/Co. However, at such potential region no catalytic\nadsorption and reaction were observed. Therefore, more interpretations are required for the difference\nof redox peaks.\nReply: We thank the reviewer for raising this important point regarding Fig. R11 (depicted as Fig.\nS23 in the revised Supporting Information). The peaks observed between –0.35 and –0.8 V vs. RHE\ndo not correspond to adsorption/catalytic events; rather, they arise from electrolyte-dependent\nreductive stripping of surface Cu/Co oxides/hydroxides and specifically adsorbed OH⁻/anion adlayers\n(e.g., NO ⁻), corresponding to transitions such as Cu O/Cu(OH) → Cu0 and Co(OH) /Co O →\n2 2 2 2 3 4\nCo0.53,90 These conditioning/oxide-stripping processes precede the catalytic window. Consistently, our\nin situ SR-FTIR detects catalytic intermediates (*CO/*COOH and *NH /*NCO) only at more negative\n2\npotentials (≤–0.9 V vs. RHE).\nTherefore, we have revised the caption in Fig. R11 to clarify that these redox features primarily arise\nfrom reduction/stripping behavior of Cu and Co surface states under different electrolyte conditions,\nrather than from direct catalytic adsorption of reaction intermediates.\n1 1 2 2\nelectroreduction of CO and NO ⁻. (A-B) CV for HER, CO RR, NO ⁻RR, and electroreduction of CO and NO ⁻ within\n2 2 2 2 2 2\nthe potential range of (A) 0.85 to -1.4 V vs. RHE and (B) 0.85 to -0.1 V vs. RHE. In panel (B), red triangles indicate oxide\nformations, while black triangles represent oxide reductions.91-94 During HER and CO RR, a single redox pair is observed,\n2\ncorresponding to the intrinsic redox transitions of the catalyst surface. In contrast, during NO ⁻RR and electroreduction of\n2 2 2\nCVs of panel (A). During HER, the cathodic feature labeled (1) at ~-0.35 V vs. RHE is assigned to reductive stripping of\ncobalt surface oxides/hydroxides (e.g., Co(OH) /Co O → Co0, consistent with the analogous peak on the Co electrode in\n2 3 4\nFig. R12B.82,96 Similar cathodic features (peak 2-5) appear under CO RR, NO ⁻RR, and electroreduction CO and NO ⁻.\n2 2 2 2\nThese are attributed to reductive of residual stripping of Cu/Co surface oxides/hydroxides and desorption of specifically\nadsorbed anion/OH⁻ adlayer accumulated during open circuit potential. Representative steps include: Cu O/Cu(OH) →\n2 2\nCu0 and Co(OH) /Co O → Co0).53,90 These conditioning/oxide-stripping peaks occur between ~-0.35 and -0.8 V vs.\n2 3 4\nRHE—preceding the catalytic window (more negative potentials) where steady CO RR/NO ⁻RR currents develop and in\n2 2\nsitu SR-FTIR detects *CO/*COOH and *NH /*NCO intermediates.\n2 2 2 3 2\n0.2 M KHCO solution, Ar-saturated 0.2 M KHCO + 0.02 M KNO solution, and CO -saturated 0.2 M KHCO + 0.02 M\n3 3 2 2 3\nKNO solution, respectively.\n2\nobserved. On the contrary, both Cu and CuCo show reversible or quasi-reversible redox peaks. The\nrelated explanations were absent in the current manuscript.\nReply: We thank the reviewer for highlighting the differences in redox behavior observed in Fig. R2.\nIndeed, the CV measurements show that pure Co displays largely irreversible redox features, while Cu\nexhibits reversible redox peaks. In contrast, the Cu–Co tandem catalyst shows quasi-reversible\nbehavior. We interpret this as follows: metallic Co undergoes surface oxidation (Co0 → Co2+/Co3+),\nbut the resulting Co2+/Co3+ species are prone to partial dissolution into the electrolyte, leading to\nirreversibility.62,63 Incorporating Cu mitigates this instability by forming a stable Cu–Co interface that\nsuppresses Co dissolution. This interfacial interaction alters the electronic structure of Co, improves\nreversibility in the redox transitions, and prevents aggregation of individual metal atoms, thereby\nenhancing overall catalyst stability.58-60 As a result, the Cu–Co system exhibits quasi-reversible\nbehavior, enabling more stable cycling with minimal loss of activity.\nThis stabilization effect is corroborated by electrochemical impedance spectroscopy (EIS), which\nshows that Cu₁Co₁ exhibits markedly lower charge-transfer resistance compared to either pure Cu or\npure Co (Fig. R3, depicted as Fig. S27 in the revised Supporting Information). The reduced resistance\nreflects improved interfacial charge transport arising from synergistic electronic interactions between\nCu and Co, which not only facilitate efficient CO and NO ⁻ but also underpins the enhanced C–N\n2 2\ncoupling performance of the bimetallic catalyst.\nTo improve clarity, we have added this explanation to the main manuscript (Page 8, Lines 140-144 &\nPage 11-12, Lines 237-247). We believe this integration provides a more comprehensive\nunderstanding of the CV redox features and their relevance to the active state of the catalysts under\noperating conditions.\nPage 8, Lines 140-144\n“Notably, Cu Co exhibits quasi-reversible redox behavior, in contrast to the reversible transition of\n1 1\npure Cu and the irreversible features of pure Co (Fig. S24).58-60 The irreversible nature of pure Co is\nlikely linked to Co2+/Co3+ dissolution into the electrolyte.62,63 Together, these findings suggest that\nCu–Co bimetallic occupy an intermediate regime, likely due to electronic and structural stabilization\nat the interface.97”\nPage 11-12, Lines 237-247\n“Electrochemical CV measurements further highlight this synergy, as the Cu–Co bimetallic catalyst\nexhibits intermediate quasi-reversible behavior between pure Cu and Co (Fig. S24). This suggests\nthat Cu suppresses Co dissolution by stabilizing its redox transitions, while Co mitigates Cu over-\noxidation.58-60 Such mutual stabilization likely arises from electronic and structural interactions at the\nCu–Co interface, further reinforcing the unique cooperative role of the tandem system. Consistently,\nelectrochemical impedance spectroscopy (EIS) reveals that Cu Co has a lower charge transfer\n1 1\nresistance compared to pure Cu and Co (Fig. S27), underscoring its enhanced electronic synergy that\n28\nfacilitates CO and NO ⁻ and underpins the C–N coupling performance of the bimetallic catalyst.\n2 2\nImportantly, this effect cannot be replicated by simple physical mixing: a drop-cast Cu/Co composite\nproduced only ammonium, formate, H , CO, and C alcohols (Fig. S33), with no detectable urea.\n2 2+\nThis finding demonstrates that atomic-level proximity achieved by co-sputtering is critical for\ncooperative interactions and the relay-type mechanism required for urea synthesis.”\ndifferent cycling numbers instead. Please correct the figure labels and descriptions accordingly.\nReply: We thank the Reviewer for pointing out this discrepancy. We have corrected the labels and\ndescriptions of Fig. R12 (adapted as Figs. S25 in the revised Supporting Information) to accurately\nreflect that the data corresponds to different cycling numbers rather than scan rates. The revised figures\nand captions in the Supporting Information now clearly indicate the cycling numbers used in each\nexperiment.\nMoreover, to reduce redundancy, we have consolidated the two graphs into a single figure (Fig. R12),\nas follows:\nFigure R12 | Close up CV curves of Cu Co catalyst (0.85 to ‒0.1 V vs. RHE) recorded at varying cycling number\n1 1\nduring HER, CO RR, NO ⁻RR, and electroreduction of CO and NO ⁻. Enlarged CVs for (A) HER, (B) CO RR, (C)\n2 2 2 2 2\nNO ⁻RR, and (D) electroreduction of CO and NO ⁻, all collected at a scan rate 20 mV s⁻1 over the potential range of 0.85\n2 2 2\nto -0.1 V vs. RHE.\nThe redox peaks of the Cu Co catalyst persist across multiple cycles. With increasing cycle number, the cathodic peak\n1 1\nbecomes more pronounced, particularly under HER and NO ⁻RR. In contrast, the introduction of CO (CO RR and\n2 2 2\nelectroreduction of CO + NO ⁻) leads to a progressive decrease in peak intensity with successive scans, likely due to\n2 2\ninteractions with the CO -saturated electrolyte.98,99\n2 2 2 2 2\nKHCO solution (pH 7), Ar-saturated 0.2 M KHCO + 0.02 M KNO solution (pH 8), and CO -saturated 0.2 M KHCO +\n3 3 2 2 3\n0.02 M KNO solution (pH 7), respectively. The CV tests for CO RR, NO ⁻RR, and electroreduction of CO & NO ⁻ were\n2 2 2 2 2\nperformed in the CO -saturated 0.2 M KHCO solution (pH = 6.8), Ar-saturated 0.2 M KHCO + 0.02 M KNO solution\n2 3 3 2\n(pH = 8.3), and CO -saturated 0.2 M KHCO + 0.02 M KNO solution (pH = 6.8), respectively.\n2 3 2\n2 3 4\nstill generated in this condition?\nReply: We are grateful for the Reviewer’s suggestion. We believe there may have been a\nmisunderstanding. The experiments in Fig. R13 (adapted as Fig. S36A in the revised Supporting\nInformation) still involved the co-electroreduction of CO and NO ⁻ in the cathode chamber; therefore,\n2 2\nammonium is still detected. The difference lies in the anolyte composition in the anode chamber. In\nFig. R13, the anolyte is the same base electrolyte as the catholyte but without the addition of NO ⁻.\n2\npresent.\nTo avoid any potential misunderstanding, we have revised the figure caption and description in the\nSupporting Information to explicitly clarify that the anolyte contains the same base electrolyte as the\ncatholyte but without NO ⁻. This change ensures that readers can clearly distinguish between the\n2\ncatholyte and anolyte compositions and better understand the experimental configuration shown in\nFig. R13.\n1 1 2\n& NO ⁻ electroreduction. The overall FE at CO -saturated catholyte: (A) alkaline (KOH, pH 10 & 13) and acidic (0.2 M\n2 2\nK SO , pH 3 & 5) conditions. Anolyte solution is the same base solution as catholyte, but without NO ⁻ present.\n2 4 2\nComparison of overall FEs using CO -saturated catholyte: (B) 0.2 M K SO (pH 5) + 0.02 M KNO and (C) 0.1 M KOH\n2 2 4 2\n(pH 13) + 0.02 M KNO at different anolyte conditions. The anolytes for panel (B) are 0.2 M K SO (pH 5) and 0.2 M\n2 2 4\nKHCO , while for panel (C) are 0.1 M KOH (pH 13) and 0.2 M KHCO . The acidic electrolyte was adjusted by the addition\n3 3\nof H SO . No C‒N products are detected in both alkaline and acidic catholytes. The alkaline catholyte suppresses HER\n2 4\ngeneration and mainly generates ammonium and formate.100 The acidic catholyte produces higher FE of H . Under both\n2\nconditions, the C product, ethanol, is consistently observed. However, the C product, n-propanol, is exclusively detected\n2 3\nunder acidic conditions. This underscore that proton (H+) availability is essential in facilitating the stabilization of reaction\nintermediates and subsequent C‒C coupling for C products.100-103 All experiments were conducted at -1.2 V vs. RHE.\n2+\nNafion membrane was utilized for conditions with pH ≤ 8, while anion exchange membrane (AEM) for alkaline conditions.\nThe error bars represent the standard deviation for at least three independent measurements.\nNote: H = hydrogen, CO = carbon monoxide, CH = methane, NH ⁺ = ammonium, HCOO⁻ = formate, CH COO⁻ = acetate,\n2 4 4 3\nCH OH = methanol, C H OH = ethanol, and C H OH = propanol.\n3 2 5 3 7\nallow comparison with CuCo samples.\nReply: We thank the Reviewer for the suggestion. As recommended, we have included the linear\ncombination analysis of XANES and EXAFS for pure Cu in Fig. R14 (adapted as Fig. S51 in the\nrevised Supporting Information). Additionally, we have added in situ XAS characterization for pure\nCo in Fig. R15 (adapted as Fig.S55 in the revised Supporting Information) and the corresponding\nlinear combination fitting of XANES and EXAFS in Fig. R16 (adapted as Fig.S58 in the revised\nSupporting Information).\nFor the pure Cu, linear combination fitting analysis of in situ XAS data indicates minor fraction of\nCu O and Cu(OH) phases alongside metallic Cu. However, because linear combination fitting reflects\n2 2\nbulk-averaged phase fractions, it cannot resolve whether these oxides are surface-localized,\ncatalytically accessible, or stable under reaction conditions.46,47 We therefore regard these values as\nindicative rather than definitive. Complementary in situ synchrotron infrared (SR-FTIR)\nmeasurements further reveal that the electroreduction of CO and NO ⁻ on pure Cu is dominated by\n2 x\nCO RR, showing intense *CO/*COOH bands but only weak *CO (carbonyl) and *CN signals,\n2\ncharacteristic of urea (Fig. R1, adapted as Fig. S42 in the revised Supporting Information). These\nresults suggest that while trace Cu oxides/hydroxides exist, their coverage and lifetime are insufficient\nto stabilize polar intermediates (*NCO, *NH CO) required for C–N coupling. Consistently, no\n2\nmeasurable urea was detected by 1H-NMR or GC-MS analysis, confirming that the catalyst behaves\nessentially as metallic Cu.\nFor the pure Co, linear combination fitting analysis of in situ XAS data identifies contribution from\nCo O and Co(OH) (Fig. R16). Yet, as with the Cu system, these oxide/hydroxide species appear only\n3 4 2\nin trace amounts. Their limited surface coverage and/or transient stability under cathodic conditions—\nas further suggested by CV measurements (Fig. R4)—renders them insufficient to promote effective\nC–N coupling.\nA detailed discussion of the Cu and Co K-edge XAS results, including their catalytic implications, is\nprovided in the main manuscript (Page 16-17, Lines 357-398).\n“The results from in situ Raman are corroborated by in situ XAS analysis at Cu and Co K-edges. The\nin situ XANES of Cu Co catalyst at Cu K-edge (Fig. 5D & Fig. S44) reveals a characteristic XANES\n1 1\nconsistent with Cu(0) and a 3d10 electron configuration.25,104 A slight shift to lower effective nuclear\ncharge in the rising edge between 8980 and 8990 eV was observed under increasingly applied\npotentials, suggesting some surface copper oxides is converted to Cu(0).105 Additionally, the main peak\n(white line) during in situ conditions is shifted to lower energies (Fig. S44), consistent with a transient\n42\nreduction associated with adsorbed reactants during electrocatalytic reactions.106 The k3-weighted\nEXAFS (Fig. 5E) and normalized FT-EXAFS (Fig. 5F) at Cu K-edge confirm that the majority Cu\nretains characteristics of metallic Cu under negative potentials, indicating that Cu returns to its\nmetallic state. OCP measurements before and after reaction show that Cu Co exhibits Cu–O and Cu–\n1 1\nCu interactions at apparent distances of ~1.4 and ~2.2 Å, respectively. Additionally, LCF of XANES\nand EXAFS spectra (Fig. S49-S50) confirms that Cu(0) is the dominant component under cathodic\npotentials, with copper oxides and hydroxides (i.e., Cu(OH) ) each accounting for <10%. These minor\n2\noxides/hydroxides species re-emerge once the potential is removed, driven by re-oxidation upon\ncontact with the electrolyte.107,108 Similar interaction patterns are also observed during the CO RR\n2\nand NO ⁻RR using Cu Co and Cu (Fig. S45-S51).\n2 1 1\nIn situ XAS measurements at the Co K-edge, examining both near-edge and extended-edge structures,\nreveal a gradual decline in the oxidation state of the catalyst to metallic Co with increasing cathodic\npotential (Fig. 5G-I & Fig. S52). This change is marked by a subtle shift towards a lower effective\nnuclear charge between 7710 and 7718 eV, accompanied by a decrease in white-line intensity, both\nindicative of a reduction in the oxidation state of Co (Fig. 5G & Fig. S52). Unlike copper, which tends\nto form oxides under OCP condition, the k3-weighted EXAFS (Fig. 5H) and normalized FT-EXAFS\n(Fig. 5I) at Co K-edge demonstrate that Co predominantly retains the properties of metallic Co under\nboth OCP conditions and negative potentials. This is likely attributed to the dissolution of Co2⁺/Co3+\nspecies into the catholyte, rendering them undetectable by XAS. The FT-EXAFS analysis further shows\na consistent increase in the intensity of Co–Co coordination at an apparent distance of 2.1 Å with the\napplication of negative potential (Fig. S52). Comparable interaction patterns are also observed for\nboth CO RR and NO ⁻RR using Cu Co and Co (Fig. S53-S55). LCF of the XANES and EXAFS spectra\n2 2 1 1\nfurther verifies that metallic Co is the principal phase under both OCP and negative potential, with\ntrace of Co O and Co(OH) during cathodic potential (Fig. S56-S58).\n3 4 2\nTogether, these results establish a clear structure-activity correlation. Minor copper\noxides/hydroxides detected by LCF EXAFS/XANES may, in principle, stabilize polar intermediates\nsuch as *NCO.53-55 Similarly, cobalt oxides/hydroxides have been reported to enhance NO ⁻ activation\n2\nand hydrogenation.30,56 However, these species are present only in trace amounts, and their limited\nsurface coverage or transient lifetime appears insufficient under electroreduction conditions, as\ndepicted in CV measurements (Fig. S23). This is consistent with the in situ SR-FTIR spectra of pure\nCu, which show only weak *CN features (Fig. S42). It should also be noted that LCF provides bulk-\naveraged phase fractions and cannot resolve whether these minority species are surface-localized,\ncatalytically accessible, or stable under bias; hence, their role is considered indicative rather than\ndefinitive for the active surface. Overall, the evidence indicates that Cu alone enhances CO\n2\nadsorption and stabilizes *CO/*COOH intermediates but lacks sufficient *NH coverage; Co alone\n2\nreduces NO ⁻ to *NH /*NH but provides inadequate *CO (Fig. S42). In contrast, at the Cu–Co\n2 2 3\ninterface, CO -derived *CO/*COOH from Cu and NO ⁻-derived *NH from Co are generated in close\n2 2 2\nproximity, which greatly enhances the probability of C–N bond formation. Guided by these\nmechanistic insights, we then performed DFT modeling of the reaction pathways.”\nXANES and (B) EXAFS spectra at Cu K-edge during electroreduction CO & NO ⁻. The fittings indicate a shift towards\n2 2\nCu metals when negative applied potentials are applied.\n2 2\nin situ XAS spectra recorded at Co K-edge for Co catalyst during electroreduction of CO and NO ⁻: (A-B) The XANES\n2 2\nspectra between 7690 – 7800 eV. (C) The k3-weighted EXAFS spectra. (D) The FT-EXAFS spectra. The XANES and\nEXAFS spectra reveal similar characteristics to those of the Co reference samples. All experiments were carried out in Ar-\nsaturated 0.2 M KHCO + 0.02 M KNO solution.\n3 2\nXANES and (B) EXAFS spectra at Co K-edge during electroreduction CO & NO ⁻. The fittings indicate a shift towards\n2 2\nCo metals when negative applied potentials are applied.\nComment: The author investigated the mechanism and reaction performance of the carbon-nitrogen\ncoupling process by changing the metal ratio, conducting in-situ characterization, and performing\ntheoretical calculations. However, the discussion in the article is still not profound enough. I reject to\naccept this work in this version for publication. The main points are as follow:\nReply: We thank the Reviewer for the critical feedback regarding the depth of our discussion. We\nfully acknowledge the importance of providing a comprehensive mechanistic understanding, and we\nappreciate the suggestion to strengthen this aspect of the manuscript.\nIn response, we have substantially revised the discussion section to offer a more detailed analysis of\nthe carbon–nitrogen coupling mechanism. Specifically, we have:\n1. Expanded interpretation of experimental results – We now provide a detailed explanation\nof why the Cu–Co tandem catalyst achieves the highest urea yield at −1.2 V vs. RHE, despite\nthe competing hydrogen evolution reaction (HER) observed during CO RR at the same\n2\npotential. We also elaborate on the rationale behind the choice NO ⁻/NO ⁻ as nitrogen sources.\n3 2\n2. Correlation of in situ spectroscopic insights – We have expanded the discussion of in situ\nsynchrotron XAS results and correlated them with in situ synchrotron based Fourier-transform\ninfrared (SR-FTIR) analysis and complementary analyses. Particular emphasis has been placed\non the dynamic changes in oxidation states and coordination environments, and their\nmechanistic relevance to C–N bond formation.\n3. Additional evidence from in situ Raman – We have included extended information on in situ\nRaman measurements, particularly regarding the identification of urea-related peaks, to further\nsubstantiate our findings.\nWe trust that the revisions presented herein adequately resolve the Reviewer’s concerns while\nsubstantially strengthening the mechanistic discussion of our work.\nshow the maximum urea performance in this point. Some reasonable explanations need to be provided.\nReply: We appreciate the reviewer for the observation. The high faradaic efficiency (FE) of HER\n(~50%) at −1.2 V vs. RHE in the CO\nreflects CO mass-transport limitation in neutral KHCO at high overpotential. As applied potential\n2 3\nbecomes more negative, the consumption rate of CO at the electrode surface can exceed its diffusion\n2\nfrom bulk electrolyte.109,110 Once interfacial CO concentration is depleted, the Volmer step dominates\n2\non Cu, so HER rises sharply.110 By contrast, when both CO and NO ⁻ are present, two coupled effects\n2 2\nsuppress HER and favor C–N coupling at the same potential:\n1. Tandem kinetics at the Cu–Co perimeter.95 At −1.2 V vs. RHE, pure Cu supplies a high\nCO/COOH flux, while pure Co rapidly reduces NO ⁻ to NH /NH (Fig. R4A, depicted as Fig.\n2 2 3\nS33A in the revised Supporting Information). Consequently, combining CO + NO ⁻ facilitates\n2 2\nthe C–N coupling with highest rate near −1.2 V vs. RHE, as shown in Fig. R17. Consistently,\nin situ SR-FTIR on Cu Co shows *CO/*COOH and *NH bands during CO RR and NO ⁻RR,\n1 1 2 2 2\nrespectively (Fig. R1, depicted as Fig. S42 in the revised Supporting Information). Urea’s\ncarbonyl (*CO) and *CN bands are clearly observed between −1 V to −1.3 V vs. RHE.\n2. HER competition and site blocking. On Cu Co , NO ⁻ serves a preferred electron/proton\n1 1 2\nsink, occupies H-binding sites, lowers *H coverage and thereby kinetically outcompetes\nHER.95 This is reflected in the CV measurements of Cu–Co tandem (Fig. R11, depicted as Fig.\nS23 in the revised Supporting Information). During HER and CO RR, one pair redox was\n2\nobserved most likely due to intrinsic redox of metal surface. However, when CV measurements\nconducted in NO ⁻RR or CO + NO ⁻, additional redox pair emerged. This might be due to\n2 2 2\ninteraction of NO ⁻ with the catalyst surface.111\neven\nthough HER is substantial during CO RR at the same potential. Taken together, these observations\n2\nsupport the feasibility of achieving high urea production rates at −1.2 V vs. RHE. Accordingly, we\nhave revised the discussion in the main manuscript (Page 9, Line 176-182) as follows:\n“Although CO RR exhibits high H formation at this potential (Fig. 3C), the mixed CO + NO ⁻ feed\n2 2 2 2\nsuppresses HER through two effects: (i) NO ⁻-driven electron–proton sinking and site blocking—NO ⁻\n2 2\nmight occupy H-binding sites and lower HER.95,112 This is consistent with the additional redox couple\nthat appears in CV curves for Cu Co upon NO ⁻ addition (Fig. S23-S25), indicative of NO ⁻-metal\n1 1 2 2\ninteraction on the catalyst surface.111 (ii) Tandem perimeter kinetics of Cu‒Co that co-generate\n*CO/*COOH and *NH in close proximity (Fig. S33), as confirmed through in situ SR-FTIR (Fig.\n2\n1 1 1 1 2\nNO ⁻RR, and electroreduction of CO and NO ⁻ (scan rate 20 mV s⁻1, 5th scan shown). Currents are normalized to the\n2 2 2\ngeometric surface area of the electrode. (B) The FE and yield rate from the NO ⁻RR. (C) The FE of the CO RR products.\nNH4⁺ 2 2\n(D) The FE and yield rate from the electroreduction of CO and NO ⁻. (E) Representative 1H NMR spectrum\nCO(NH2)2 2 2\nshowing CO(NH ) produced at ˗1.2 V vs. RHE. (F) Urea detection by 1H-NMR at t = 0, 5, and 12 h of the experiment at\n2 2\n˗1.2 V vs. RHE. The CO RR, NO ⁻RR, and electroreduction of CO and NO ⁻ were carried out in CO -saturated 0.2 M\n2 2 2 2 2\nKHCO , Ar-saturated 0.2 M KHCO and 0.02 M KNO , and CO -saturated 0.2 M KHCO and 0.02 M KNO solutions,\n3 3 2 2 3 2\nrespectively.\noxidation state of copper were analyzed, without discussing the key steps of the carbon-nitrogen\ncoupling process from the perspective of the dynamic changes of the intermediates.\nReply: We thank the Reviewer for this comment. We respectfully note that in addition to analyzing\nthe oxidation states of Cu and Co, we have substantially expanded our discussion to include the\ndynamic evolution of surface intermediates probed by in situ Raman, SR-FTIR, and isotope labeling.\nSpecifically:\n In situ Raman detected the appearance of NO-derived (~730 cm⁻1) and CO-derived (~2116\ncm⁻1) intermediates, followed by urea-related vibrations at ~1002 cm⁻1 (C‒N) and ~575 cm⁻1,\nwhich were absent in the CO RR and NO ⁻RR (Fig. R18A, depicted as Fig. 5A in the revised\n2 2\nmanuscript).\n In situ SR-FTIR revealed *NH 2 (~3115 cm⁻1), *CO (2065 cm⁻1), and the first coupling\nintermediate *NCO (~1694 cm⁻1) vibrations (Fig. R18B–C, depicted as Fig. 5B-C in the\nrevised manuscript). These were subsequently followed by urea-specific *CO (carbonyl, ~1590\ncm⁻1) and *CN (~1455 cm⁻1) that match the urea reference (Fig. R19, depicted as Fig. S40 in\nthe revised Supporting Information). Notably, these urea-specific vibrations (*CO carbonyl\nand *CN) were absent during CO RR & NO ⁻RR (Fig. R20, depicted as Fig. S41 in the revised\n2 2\nSupporting Information), as well as in the monometallic catalyst controls (R1, depicted as Fig.\nS41 in the revised Supporting Information).\n 15N isotopic substitution further confirmed that the observed *CN modes originate from NO x ⁻-\nderived nitrogen (Fig. R9, depicted as Fig. S43 in the revised Supporting Information).\n These intermediate evolutions were correlated with in situ XAS, which demonstrates that Cu\npredominantly stabilizes *CO/*COOH, Co promotes *NH /*NH formation, and only the Cu–\n2 3\nCo interface facilitates their co-adsorption and subsequent C–N bond formation.\nTogether, these measurements provide a mechanistic picture of C–N coupling that goes beyond\nchanges in metal oxidation states, directly addressing the Reviewer’s request to discuss the key steps\nfrom the perspective of intermediates.\nThe updated discussion that includes elaborates change in catalyst structure and intermediate formation\ncan be found in main manuscript (Page 14-17, Line 317-398):\n“To correlate surface intermediates with catalyst structure, we combined in situ Raman and SR-FTIR,\nand XAS. In situ Raman spectroscopy (900 to 3400 cm⁻1) tracked the evolution of intermediates during\nCO and NO ⁻ electroreduction (Fig. 5A & Fig. S39). Under open circuit potential (OCP), bands from\n2 2\nbidentate nitrite (νNO , ~1360 cm⁻1) and buffer solution (HCO ⁻, ~1062 cm⁻1) were observed and\n2 3\npersisted at −1.2 V vs. RHE.113-115 After 20 minutes at −1.2 V vs. RHE, new features at ~730 (NO\nbending) and ~2116 cm⁻1 (CO stretching) appeared, consistent with the reduction of NO ⁻ and CO ,\n2 2\nrespectively.116-118 Urea related vibrations at 1002 cm⁻1 (C‒N) and 575 cm⁻1 (amide OCN) emerged\nafter 30 minutes, but were absent in the control CO RR or NO ⁻RR (Fig. S38), confirming its specific\n2 2\n41\nassociation with C–N coupling.116 Similar intermediates were identified under different applied\npotentials (Fig.S39), with C‒N bond and OCN bending identified from −0.9 V vs. RHE. Nevertheless,\nthe amine group intermediate signals (e.g.: *NH ) resulted from further reduction of NO ⁻ and NO are\n2 2\nlikely undetectable in Raman spectroscopy due to weaker NH vibrations and low laser power.119\nat various applied potentials (Fig. 5B,C). Bands assigned to *NH bending mode (~3115 cm⁻1), *CO\n2\nstretching (~2065 cm⁻1), *NH rocking vibration (~1150 cm⁻1) and *NH rocking/wagging mode (~930\n2 , 3\ncm⁻1) appeared from −0.8 V vs. RHE, indicative of CO and NO ⁻ reduction.14,44,50,87,120 Notably, an\n2 2\namide-related *NCO vibrations (e.g. *NHCO, *NH CO, or *NH COOH), emerged around 1694 cm⁻1\n2 2\nfrom −0.9 V vs. RHE, suggesting first C‒N coupling formation.14,16,58 Distinct vibrational features\ncorresponding to the carbonyl (*CO) and *CN bonds of urea were observed at approximately 1590\ncm⁻1 and 1455 cm⁻1, respectively, over the potential range of −0.9 V to −1.3 V vs. RHE, closely\nresembling the C=O and C‒N stretching frequencies of the urea reference (1589 cm⁻1 and 1457 cm⁻1,\nrespectively; Fig. S40). In contrast, control CO RR and NO ⁻RR experiments display only\n2 2\n*CO/*COOH and *NH /*NH features (Fig. S41), without any C–N features, thereby implicating\n2 3\n*CO/*COOH and *NH as precursors for C‒N coupling.14,16,17,58\nand *CN vibrational features were exclusively observed for Cu Co (Fig. S42). Pure Cu exhibited\n1 1\nmainly *CO (~2065 cm⁻1) and *COOH (~1720 cm⁻1) bands associated with gaseous CO and formate\nformation, along with a C‒O stretching band (*CO) at ~1030 cm⁻1 corresponding to methanol.52 These\nresults indicate that CO RR dominates on pure Cu, consistent with earlier observations. In contrast,\n2\npure Co displayed no carbonyl or *CN bands, yielding mainly formate and ammonia. These results\nindicate that the Cu–Co interface uniquely promotes the C–N coupling steps leading to urea.\nIsotope labelling with 15NO ⁻ confirmed the nitrogen origin of the *CN band (Fig. S43). Spectra for\n3\n14NO ⁻ and 14NO ⁻ were nearly identical, apart from an additional *NO ⁻ feature at ~1340 cm⁻1.87\n2 3 3\nSubstitution with 15NO ⁻ shifted nitrogen-containing vibrational modes to lower wavenumbers, with\n3\nthe *C15N feature of urea detected at 1411 cm⁻1, while the carbonyl peak remained unchanged.88 This\nisotopic shift conclusively verifies that the *CN signal originates from NO ⁻-derived nitrogen.\nx\nCatalyst structure and dynamics were probed by in situ Raman and XAS. In situ Raman spectra in the\n300-1000 cm⁻1 region showed no oxides or hydroxides signals under negative potentials (Fig. S39),\nsuggesting that the active sites during the electroreduction are in the metallic forms of Cu and Co.\nOnly at OCP, surface hydroxides (Co(OH) and Cu(OH) ) were detected at ~502.5 cm⁻1, implying the\n2 2\ninteraction between Cu/Co and the electrolyte.30,80,121\nThe results from in situ Raman are corroborated by in situ XAS analysis at Cu and Co K-edges. The\nin situ XANES of Cu Co catalyst at Cu K-edge (Fig. 5D & Fig. S44) reveals a characteristic XANES\n1 1\nconsistent with Cu(0) and a 3d10 electron configuration.30,80,121 A slight shift to lower effective nuclear\ncharge in the rising edge between 8980 and 8990 eV was observed under increasingly applied\npotentials, suggesting some surface copper oxides is converted to Cu(0).105 Additionally, the main peak\n(white line) during in situ conditions is shifted to lower energies (Fig. S44), consistent with a transient\n42\nreduction associated with adsorbed reactants during electrocatalytic reactions.106 The k3-weighted\nEXAFS (Fig. 5E) and normalized FT-EXAFS (Fig. 5F) at Cu K-edge confirm that the majority Cu\nretains characteristics of metallic Cu under negative potentials, indicating that Cu returns to its\nmetallic state. OCP measurements before and after reaction show that Cu Co exhibits Cu–O and Cu–\n1 1\nCu interactions at apparent distances of ~1.4 and ~2.2 Å, respectively. Additionally, LCF of XANES\nand EXAFS spectra (Fig. S49-S50) confirms that Cu(0) is the dominant component under cathodic\npotentials, with copper oxides and hydroxides (i.e., Cu(OH) ) each accounting for <10%. These minor\n2\noxides/hydroxides species re-emerge once the potential is removed, driven by re-oxidation upon\ncontact with the electrolyte.107,108 Similar interaction patterns are also observed during the CO RR\n2\nand NO ⁻RR using Cu Co and Cu (Fig. S45-S51).\n2 1 1\nIn situ XAS measurements at the Co K-edge, examining both near-edge and extended-edge structures,\nreveal a gradual decline in the oxidation state of the catalyst to metallic Co with increasing cathodic\npotential (Fig. 5G-I & Fig. S52). This change is marked by a subtle shift towards a lower effective\nnuclear charge between 7710 and 7718 eV, accompanied by a decrease in white-line intensity, both\nindicative of a reduction in the oxidation state of Co (Fig. 5G & Fig. S52). Unlike copper, which tends\nto form oxides under OCP condition, the k3-weighted EXAFS (Fig. 5H) and normalized FT-EXAFS\n(Fig. 5I) at Co K-edge demonstrate that Co predominantly retains the properties of metallic Co under\nboth OCP conditions and negative potentials. This is likely attributed to the dissolution of Co2⁺/Co3+\nspecies into the catholyte, rendering them undetectable by XAS. The FT-EXAFS analysis further shows\na consistent increase in the intensity of Co–Co coordination at an apparent distance of 2.1 Å with the\napplication of negative potential (Fig. S52). Comparable interaction patterns are also observed for\nboth CO RR and NO ⁻RR using Cu Co and Co (Fig. S53-S55). LCF of the XANES and EXAFS spectra\n2 2 1 1\nfurther verifies that metallic Co is the principal phase under both OCP and negative potential, with\ntrace of Co O and Co(OH) during cathodic potential (Fig. S56-S58).\n3 4 2\nTogether, these results establish a clear structure-activity correlation. Minor copper\noxides/hydroxides detected by LCF EXAFS/XANES may, in principle, stabilize polar intermediates\nsuch as *NCO.53-55 Similarly, cobalt oxides/hydroxides have been reported to enhance NO ⁻ activation\n2\nand hydrogenation.30,56 However, these species are present only in trace amounts, and their limited\nsurface coverage or transient lifetime appears insufficient under electroreduction conditions, as\ndepicted in CV measurements (Fig. S23). This is consistent with the in situ SR-FTIR spectra of pure\nCu, which show only weak *CN features (Fig. S42). It should also be noted that LCF provides bulk-\naveraged phase fractions and cannot resolve whether these minority species are surface-localized,\ncatalytically accessible, or stable under bias; hence, their role is considered indicative rather than\ndefinitive for the active surface. Overall, the evidence indicates that Cu alone enhances CO\n2\nadsorption and stabilizes *CO/*COOH intermediates but lacks sufficient *NH coverage; Co alone\n2\nreduces NO ⁻ to *NH /*NH but provides inadequate *CO (Fig. S42). In contrast, at the Cu–Co\n2 2 3\ninterface, CO -derived *CO/*COOH from Cu and NO ⁻-derived *NH from Co are generated in close\n2 2 2\nproximity, which greatly enhances the probability of C–N bond formation. Guided by these\nmechanistic insights, we then performed DFT modeling of the reaction pathways.”\n1 1 1 1\ncatalyst at -1.2 V vs. RHE. (B-C) In situ SR-FTIR spectra of Cu Co catalyst in the range of 900 – 2200 cm⁻1 (B) and 2800\n1 1\n– 3400 cm⁻1 (C). (D-F) In situ XAS: XANES spectra (D), k3-weighted EXAFS spectra (E), and FT-EXAFS spectra (F) of\nCu Co catalyst at Cu K-edge. (G-I) In situ XANES spectra (G), k3-weighted EXAFS spectra (H), and FT-EXAFS spectra\n1 1\n(I) of Cu Co catalyst at Co K-edge. All experiments were conducted in a CO -saturated 0.2 M KHCO and 0.02 M KNO\n1 1 2 3 2\nsolution.\nFTIR of urea shows peak approximately at 1150, 1457, 1589, and 1672 cm⁻1 attributed to rocking vibration of NH ,\n2\nstretching frequency of C‒N and C=O bonds, and bending frequency of NH . The symmetric and asymmetric vibration of\n2\n1 1\nFTIR) spectra during (A) CO RR, (B) NO ⁻RR, and (C) electroreduction of CO and NO ⁻. The CO RR, NO ⁻RR, and\n2 2 2 2 2 2\nelectroreduction of CO and NO ⁻ were carried out in the CO -saturated 0.2 M KHCO solution, Ar-saturated 0.2 M KHCO\n2 2 2 3 3\nand 0.02 M KNO solution, and CO -saturated 0.2 M KHCO and 0.02 M KNO solution, respectively. The *COOH and\n2 2 3 2\n*CO peaks related to formate and CO gas were observed at approximately 1720 and 2065 cm⁻1, respectively, during\nCO RR.50-52 The C‒O bond related to alcohol (i.e. methanol) was also detected around 1030 cm⁻1.52 For NO ⁻RR, the peaks\n2 2\nat around 930, 1150, 1250, and 1610 cm⁻1 were attributed to *NH , *NH , *NO ⁻, and *NO, accordingly.87,120,123-126 During\n3 2 2\nco-reduction of CO + NO ⁻, the *COOH vibrational feature is not prominently observed; instead, a distinct *NCO band\n2 2\n(~1694 cm⁻1) appears. This suggests that *COOH is likely a short-lived intermediate during electrochemical C–N coupling,\nor that its signal is obscured by overlapping features from other intermediates (i.e., *NCO).\n2\ncommonly used nitrate is ineffective or inefficient in this case requires more experimental evidence to\nsupport it.\nReply: We thank the reviewer for this helpful comment. We agree that our rationale for choosing NO ⁻\n2\nover NO ⁻ should be articulated more clearly. Mechanistically, NO ⁻ reduction is gated by a slow,\n3 3\nhigh-overpotential 2e⁻ step (NO ⁻→ NO ⁻) that competes strongly with HER in neutral KHCO .127,128\n3 2 3\nThis might result in low steady-state surface coverage of NH to couple efficiently with *CO, as\n2\nevidence of low yield rate of urea when NO ⁻ is applied (Fig. R21, depicted as Fig. 4 in the revised\n3\nmanuscript). In contrast, feeding NO ⁻ bypasses this kinetic bottleneck, enabling rapid formation of\n2\n*NH on Co at the same potentials where Cu supplies CO, thereby maximizing the C–N coupling rate.\n3 2\nS22 in the revised Supporting Information): (i) potential-dependent urea FE shows with NO ⁻ urea\n3\nappears later at −1 V vs. RHE with highest yield observed at −1.4 V vs. RHE, while NO\n2\n⁻ yields urea\nfrom −0.9 V vs. RHE and pronounces urea maxima at −1.2 V vs. RHE; and (ii) in situ SR-FTIR with\nNO ⁻ shows weaker *NH band, as well as late formation of *NCO/urea bands in the catalytic window,\n3 2\nwhereas NO ⁻ produces clear *NH , *NCO and urea carbonyl/*CN features (Fig. R9, depicted as Fig.\n2 2\nS43 in the revised Supporting Information). These results are now summarized in the main manuscript\n(Page 10, Lines 190-207).\n“When NO ⁻ was used as a nitrogen source instead of NO ⁻, similar products were formed but at lower\n3 2\nfaradaic efficiencies and yield rates (Fig. 4A and Fig. S30). Urea was detected only at more negative\npotentials (≤ −1.1 V vs. RHE) and achieved maximum values of 8 ± 1% and 26 ± 4 mmol h⁻1 g ⁻1\ncat\n(0.018 ± 0.003 nmol s⁻1 cm ⁻2) at −1.3 V vs. RHE (Fig. S31). This difference is consistent with our\nECSA\nprevious findings, where the activation and reduction of NO ⁻ require lower overpotentials compared\n2\nto NO ⁻, enabling the preferential reduction of prior reactant.107 Supporting evidence was provided by\n3\nin situ SR-FTIR spectra (Fig. S42), in which characteristic urea-related vibrations (*NCO, *CO\ncarbonyl, and *CN) were observed more weakly in CO + NO ⁻ than in CO + NO ⁻ experiments.129,130\n2 3 2 2\nTo further probe nitrogen source effects, hydroxylamine (NH OH) was tested. Instead of urea, the\n2\nreaction yielded ammonium carbamate ([NH ][NH COO]) with FE of 16 ± 3% and yield rate of 90 ±\n4 2\n5 mmol h‒1 g ‒1 (0.064 ± 0.004 nmol s⁻1 cm ⁻2) at −1.2 V vs. RHE (Fig. 4A). This suggests that the\ncat ECSA\ndominant process is electrochemical reduction of NH OH to NH +, while the carbamate anion forms\n2 4\nchemically through reaction of hydroxylamine with CO .129,130 The absence of ammonium carbamate\n2\nin NO ⁻ and NO ⁻ experiments implies that hydroxylamine is not an intermediate in the co-reduction\n2 3\npathway, or if transiently formed, is rapidly converted into other species. Overall, these results\nestablish that NO ⁻ is a more effective nitrogen source than NO ⁻ or NH OH for urea electrosynthesis.\n2 3 2\nCompared with reported catalysts (Fig. 4B,C & Table S5), the Cu Co system achieves competitive\n1 1\nurea yield rates under prolonged operation, underscoring its potential as a tandem Cu–Co catalyst\nfor selective C–N coupling.”\nyield rate of C‒N products from different nitrogen sources. The concentration of nitrogen sources was kept at 0.02 M in\nthe 0.2 M KHCO solution. The applied potential for the test was ˗1.2 V vs. RHE. (B-C) Comparison of applied potential\n3\n(B) and FE (C) with urea yield rates achieved from CO and NO ⁻/NO ⁻ in this work versus those reported in the literature\n2 2 3\n(Table S5). (D) CV curves of Cu (n) Co (1-n) catalysts during electroreduction of CO 2 and NO 2 ⁻. (E) The FE of C‒N products\nusing Cu (n) Co (1-n) catalysts after electroreduction of CO 2 and NO 2 ⁻ for 12 h. (F) The electroreduction of CO 2 and NO 2 ⁻ in\nvarious catholyte and anolyte at ˗1.2 V vs. RHE. For panel (A) electroreduction of CO and NO ⁻ were carried out in CO -\n2 2 2\nsaturated 0.2 M KHCO with addition of 0.02 M KNO or 0.02 M KNO or 0.02 M NH OH. Catholyte for panels (B-E) is\n3 3 2 2\nCO -saturated 0.2 M KHCO + 0.02 M KNO . The concentrations of anolyte in panel (F) are 0.1 M KOH (pH 13), 0.2 M\n2 3 2\nK SO + H SO (pH 5), and 0.2 M KHCO (pH 8). For catholyte in panel (F), the buffer concentrations are identical to\n2 4 2 4 3\nthose of the anolyte but include 0.02 M KNO . All testing was performed using a continues CO gas flow.\n2 2\n2 x\ncurves of Cu Co under different reactions. (B-D) FE and yield rate of urea from (B) CO + NO ⁻ and (C-D) CO + NO ⁻.\n1 1 2 2 2 3\nUrea yield rates in panel (B) and (D) were normalized to electrochemical active surface area (ECSA) (Fig. S23). The\nmaximum FE and yield rate of urea from CO + NO ⁻ was approximately 8 ± 1% and 26 ± 4 mmol h⁻1 g ⁻1 (0.018 ± 0.003\n2 3 cat\nnmol s⁻1 cm ⁻2), respectively. In comparison, CO + NO ⁻ achieved higher values of approximately 11 ± 2% and 61 ±\nECSA 2 2\n6 mmol h⁻1 g ⁻1 (0.043 ± 0.005 nmol s⁻1 cm ⁻2). All measurements in panel (B-D) were conducted in CO -saturated\ncat ECSA 2\n0.2 M KHCO + 0.02 M KNO solutions over 12 h reactions.\n3 x\nexperiments? Are interferences from other nitrogen-containing species excluded?\nReply: We thank the reviewer for this comment. The ~1002 cm⁻¹ Raman band is a well-established\nsignature of urea (δ(N–C–N)/skeletal C–N stretch) and is widely used in electrochemical studies.131,132\nIn our work, this assignment is further corroborated by control experiments: in situ Raman spectra of\nCO RR and NO ⁻RR individually show no such feature, whereas the band clearly emerges only under\n2 2\nco-reduction conditions. In addition, in situ SR-FTIR measurements provide complementary evidence,\nreinforcing the assignment of this ~1002 cm⁻¹ band to urea formation.\n1. In situ Raman of CO RR and NO ⁻RR: the 1002 cm⁻¹ band is absent in both CO RR and\n2 2 2\nNO ⁻RR under identical conditions, ruling out contributions from background electrolyte\n2\nspecies. Notably, potential Raman interferences (NO ~730 cm⁻1; HCO ⁻ ~1062 cm⁻1; NO ⁻\n3 2\n~1360 cm⁻¹; CO ~2116 cm⁻1) occur at distinct wavenumbers and are observed in the CO RR\n2\nand NO ⁻RR (Fig. R23, as depicted as Fig. S39 in the revised Supporting Information).\n2\n2. In situ infrared: In situ SR-FTIR in the same potential window (−0.9 to −1.3 V vs. RHE)\nshows the urea carbonyl (~1590 cm⁻1) and C–N (~1455 cm⁻1) bands that track with the Raman\n1002 cm⁻¹ feature in potential and time (Fig. R1, depicted as Fig. S41 in the revised Supporting\nInformation).\n3. Isotopic validation: 15N isotope labelling in the in situ SR-FTIR shifts the nitrogen-containing\nmodes (e.g., *CN 1455 → 1411 cm⁻¹) while leaving the carbonyl unchanged, confirming the\nNO ⁻ origin of nitrogen in the C–N product (Fig. R9, depicted as Fig. S43 in the revised Supporting\nx\nInformation).\nTogether, these controls and in situ validations strongly support the assignment of the 1002 cm⁻1\nRaman band to urea. We have expanded the discussion accordingly in the revised manuscript (Page\n15, Lines 322–324).\n“Urea related vibrations at 1002 cm⁻1 (C‒N) and 575 cm⁻1 (amide OCN) emerged after 30 minutes,\nbut were absent in the control CO RR or NO ⁻RR (Fig. S38), confirming its specific association with\n2 2\nC–N coupling.116”\n1 1\ntime intervals at ˗1.2 V vs. RHE, (B) under different potentials for 60 minutes, and (C) comparison CO + NO ⁻, CO , and\n2 2 2\nNO ⁻ electroreduction at ˗1.2 V vs. RHE. At open circuit potentials (OCP), contact between CuO/CoO and catholyte was\n2\ntranslated as signals around 502.5 cm⁻1 that can be identified as Cu(OH) and Co(OH) .30,80,121 When negative potential\n2 2\nwas applied, no metal oxides or hydroxides are detected within the range, implying that Cu(0) and Co(0) are the active\nsites during the reaction. The vibrational peak of OCN bending mode at 575 cm⁻1 appeared with high intensity since ˗1.0\nV vs. RHE or after 30 min at ˗1.2 V vs. RHE, indicating amide formation.116 The CO and C‒N bond was detected from\n˗0.9 V vs. RHE around 2116 and 1002 cm⁻1, respectively.116,118 Notably, the C–N band was absent during CO RR and\n2 2 3 2\nis not clearly indicated in the article's evidence.\nReply: We thank the Reviewer for this thoughtful comment on whether the hydrogenation process in\nurea formation is mediated by active hydrogen species (*H) or proceeds via a proton–coupled electron\ntransfer (PCET) pathway.\nIn the revised manuscript, we have clarified the following points:\n1. Evidence from pH and membrane effects: Our systematic variation of catholyte/anolyte\npH\nand membrane type strongly suggests that proton availability and H+ transfer critically govern\nthe distribution of products by steering protonation pathways of CO and/or NO ⁻ toward\n2 2\n*CO/*COOH and *NH , respectively. In particular, urea generation is suppressed under both\n3\nhighly acidic and alkaline conditions, but is favored under neutral CO -saturated KHCO with\n2 3\nNafion membrane. These observations support a PCET-driven mechanism, in which a finely\nregulated proton flux from the electrolyte/anolyte is essential for C–N coupling, rather than a\nmechanism dominated by surface-bound active hydrogen species (*H).15,16,101,133,134\n2. Correlation with situ spectra: In situ SR-FTIR (Fig. R1) and Raman (Fig. R23)\nmeasurements reveal the co-emergence of *NH and *CO/*COOH species during\n2\nelectroreduction, consistent with PCET-mediated hydrogenation steps. In contrast, if active\nhydrogen species were dominant, stronger HER activity and broader evidence of *H-related\nspecies would be expected around 3300-3400 cm⁻1, which we do not observe under urea-\nproducing conditions.135\n3. Limitations: We acknowledge that unambiguous discrimination between active hydrogen\nspecies and PCET would require isotope-labelling (2H(D)) or kinetic isotope effect studies.136\nNevertheless, based on our current experimental evidence, we propose that PCET is the more\nplausible pathway governing hydrogenation in the C–N coupling process.\nWe have revised the discussion accordingly to highlight this mechanistic interpretation while\nacknowledging the limitations of our data. The revision can be seen in the main manuscript (Page\n12-13, Lines 264-295) as follows:\n“Beyond tuning the Cu:Co ratio, we also explored the effects of the ion-exchange membrane and pH\nof the catholyte and anolyte on the effectiveness of the C‒N coupling for the urea generation. Control\nexperiments were performed by adjusting the pH of both catholyte and anolyte to acidic, neutral, and\nalkaline conditions under continuous CO -gas flow (Table S6). No urea was detected after co-\n2\nreduction of CO and NO ⁻ at −1.2 V vs. RHE under either acidic (pH 3 & 5, K SO , Nafion membrane)\n2 2 2 4\nor alkaline conditions (pH 10 & 13, KOH, anion exchange membrane (AEM)). Under alkaline\nconditions, formate was the dominant product, acidic electrolytes suppressed C‒N coupling and\nsteered the reaction toward H and ammonium generation (Fig. S36). Urea formation was consistently\n2\nobserved only when KHCO was used as buffer in both catholyte and anolyte, together with Nafion as\n3\nmembrane (Fig. 4F & Fig. S31). Notably, replacing Nafion membrane with AEM under identical\nconditions eliminated urea production (Table S6). These results indicate that controlled proton\n52\navailability is critical for regulating surface protonation steps, consistent with a PCET pathway rather\nthan mediation by surface-bound hydrogen (*H).101,133,134 This conclusion is further supported by in\nsitu SR-FTIR measurements (Fig. S41), which show the absence of broad vibrational signatures\ncharacteristic of *H-related species in the 3300–3400 cm⁻1 region.135,137\nTo rationalize these trends, we consider the electrolyte-dependent mechanistic factors. Neutral\nKHCO maintains relatively high dissolved CO concentrations, in contrast to alkaline conditions\n3 2\nwhere CO is converted to HCO ⁻/CO 2⁻, thereby enabling robust *CO/*COOH formation.138 In\n2 3 3\nacidic electrolytes, *CO can form but HER dominates surface sites and mass transport, strongly\ninhibiting C–N coupling.139 For NO ⁻ reduction, neutral pH provides sufficient proton activity to drive\n2\n*NO hydrogenation to *NH without excessive HER, whereas limited proton availability in alkaline\n2\nelectrolytes slows hydrogenation steps and reduces *NH coverage.123,140,141 Furthermore, neutral\n3 2\nwithout severe *OH blocking that occurs in strongly alkaline media.110,142,143\nAnolyte variation further highlighted the importance of proton flux (Table S6). Maintaining the\ncatholyte at neutral KHCO but lowering the anolyte to pH 5 enabled detectable urea (FE 2 ± 0.2%,\n3\nyield rate 14 ± 2 mmol h‒1 g ‒1; Fig. 4F & Fig. S37), whereas excessive proton concentration (pH 3)\ncat\ndiminished CO availability and enhanced competing H /NH + pathways. This delicate balance in H⁺\n2 2 4\nconcentration highlights the interplay between CO activation and NO ⁻ hydrogenation, in which\n2 2\noptimal PCET-driven protonation enables co-coverage of *CO/*COOH and *NH .23,133,136,144\nvibrational bands under neutral conditions (Fig. S41-S42). Collectively, these results emphasize that\nurea formation requires not only an optimal Cu–Co surface but also finely tuned proton management\nacross both membrane and electrolyte compartments.”\nmechanism still needs to be further enhanced.\nReply: We acknowledge the Reviewer’s point that the overall urea performance remains modest\ncompared to the literature. While our catalyst does not yet achieve record-high efficiency, its\nsignificance lies in establishing a fundamental mechanistic framework for electrochemical urea\nsynthesis. By integrating systematic electrolyte/membrane control experiments with in situ\nspectroscopic investigations (XAS, SR-FTIR, Raman), we demonstrate that: (i) proton management\nand pH balance critically govern C–N coupling selectivity (see response to Comment 5); (ii) catalyst\nstability is compromised under acidic and alkaline conditions but preserved in neutral electrolytes;\nand\n(iii) tandem Cu–Co sites enable proximity-driven co-adsorption of *CO/*COOH and *NH (see\n2\nresponse to Comments 3 and 5). These experimental insights are further corroborated by DFT\nmodeling, which connects the spectroscopic evidence with feasible C–N coupling pathways.\nThus, while the absolute urea performance can certainly be improved through future catalyst design,\nthe present study makes a unique contribution by strengthening the mechanistic understanding of the\nC–N coupling process, thereby addressing a critical gap in the field.\nWe have also expanded the discussion of the role of Cu–Co tandem catalysis in electrochemical C–N\ncoupling in the main manuscript (Page 10, Lines 208–218; Page 11–12, Lines 232–247), as well as\nthe influence of pH balance on catalyst stability (Page 13, Lines 296–308).\nPage 10, Lines 208-18\n“We also investigated electroreduction of CO\n2\nand NO\n2\n– at −1.2 V vs. RHE using monometallic Cu\nand Co catalysts (Fig. S30 & Fig. S33) as control references. The sputtered Cu catalyst produced no\ndetectable C–N products, instead yielding formate, ethanol, and propanol (Fig. S29, S47-S48). This\noutcome contrasts with many reports describing strong C–N coupling on Cu and underscores a more\nnuanced picture.15,48 In the literature, high urea selectivity is typically associated with oxide-derived\nCu, where mixed Cu(0)/Cu(I) and defect-rich surfaces stabilize polar intermediates such as *CO,\n*NCO and *NH CO.42,44 By comparison, sputtered Cu typically generates thin, relatively smooth films\n2\nwith fewer defect sites, thereby limiting the stabilization of key intermediates required for C‒N\ncoupling.45 Consequently, metallic Cu favors CO adsorption and *CO/*COOH formation,49\n2\nconsistent with the results observed in the in situ SR-FTIR spectra (Fig. S42). For pure Co, the main\nproducts were ammonium and H (Fig. S33), with no evidence of C–N products, also corroborated by\n2\nin situ SR-FTIR (Fig. S42).”\nPage 11-12, Lines 232-247\n“These controls underscore that neither Cu nor Co alone is sufficient for effective C–N coupling.\nInstead, the tandem interface of Cu–Co is essential: Cu sites efficiently adsorb and activate CO to\n2\n*CO, while Co sites accelerate the stepwise reduction of NO ⁻ to *NH /*NH . The spatial co-\n2 2 3\ngeneration and proximity of *CO and *NH at the Cu–Co perimeter greatly enhances the likelihood\n2\nof C–N bond formation, leading to markedly higher urea activity and selectivity than either metal\nalone.\nexhibits intermediate quasi-reversible behavior between pure Cu and Co (Fig. S24). This suggests\nthat Cu suppresses Co dissolution by stabilizing its redox transitions, while Co mitigates Cu over-\noxidation.58-60 Such mutual stabilization likely arises from electronic and structural interactions at the\nCu–Co interface, further reinforcing the unique cooperative role of the tandem system. Consistently,\nelectrochemical impedance spectroscopy (EIS) reveals that Cu Co has a lower charge transfer\n1 1\nresistance compared to pure Cu and Co (Fig. S27), underscoring its enhanced electronic synergy that\nfacilitates CO and NO ⁻ and underpins the C–N coupling performance of the bimetallic catalyst.\n2 2\nImportantly, this effect cannot be replicated by simple physical mixing: a drop-cast Cu/Co composite\nproduced only ammonium, formate, H , CO, and C alcohols (Fig. S33), with no detectable urea.\n2 2+\nThis finding demonstrates that atomic-level proximity achieved by co-sputtering is critical for\ncooperative interactions and the relay-type mechanism required for urea synthesis.”\nPage 13, Lines 296-308\n“The role of electrolyte pH on catalyst stability is further reflected in CVs recorded in the non-catalytic\npotential window (Fig. S34). Acidic electrolytes promoted rapid Cu/Co redox cycling, consistent with\naccelerated dissolution,145,146 whereas alkaline electrolytes induced broad, hysteretic oxidation waves\ncharacteristic of irreversible oxide/hydroxide formation.143 By contrast, neutral KHCO displayed\n3\nonly weak and broadened transitions, indicating moderate surface dynamics. These CV signatures\ndemonstrate that acidic and alkaline conditions drive aggressive redox activity and oxide formation,\ncompromising structural integrity, whereas neutral conditions preserve a stable Cu–Co interface.\nThe instability was corroborated by ICP analysis of Cu and Co dissolution under strongly acidic and\nalkaline environments (Fig. S35). In acidic media (pH 3), Cu and Co dissolution achieved 8 ± 2 wt.%\nand 17 ± 2 wt.%, respectively. In alkaline media (pH 13), dissolution was lower (3 ± 1 wt.% for Cu\nand 8 ± 2 wt.% for Co). Such losses significantly undermine the long-term stability and performance\nof the catalysts. Finally, we observed that urea is prone to decomposition into ammonium and\nformate/formic acid under highly acidic (pH 3, K SO ) and alkaline (pH 13, KOH) conditions (Fig.\n2 4\nS38), especially at low concentrations.”\nComment: This manuscript presents a compelling and thorough investigation into the co-\nelectroreduction of CO and NO ⁻ to urea using co-sputtered bimetallic Cu-Co catalysts. In situ\n2 2\nsynchrotron X-ray absorption spectroscopy was employed to demonstrate dynamic stability of metallic\nCu and Co phases under operating conditions. Besides, the authors further proposed synchrotron-based\ninfrared spectroscopy to investigate the specific reaction mechanism. DFT calculations provide further\nevidence for the potential determining step. Finally, they achieve a high urea yield rate of 61 ± 6 mmol\nh‒1 g ‒1 at –1.2 V vs. RHE. The work is well-structured, given the growing interest in sustainable\ncat\nnitrogen chemistry, and offers meaningful contributions to the understanding of C-N coupling\nmechanisms. However, before it could be accepted, some questions still need to answer:\nReply: We sincerely thank the Reviewer for their positive assessment of our manuscript and for\nrecognizing the significance of our work on co-electroreduction of CO and NO ⁻ to urea. We greatly\n2 2\nappreciate the Reviewer’s comments highlighting the strengths of our study. In response, we have\nsubstantially revised the discussion section to provide a more detailed and rigorous analysis of the\ncarbon–nitrogen coupling mechanism. Specifically, we have:\n1. The atomic level-mixing of co-sputtered Cu‒Co: Performed additional experiments using\ndrop-cast Cu/Co catalysts, which did not yield urea, thereby confirming the critical role of\natomic-level mixing uniquely achieved by co-sputtering.\n2. XPS analysis: Revised and improved the deconvolution of the Co 2p spectra to ensure more\naccurate phase identification.\n3. Mechanistic rationale for neutral pH: Expanded the discussion of pH dependence in urea\nformation, integrating experimental evidence and mechanistic insights into proton-coupled\nelectron transfer (PCET).\n4. Normalization to ECSA: Normalized current densities (j) and urea yield rates to\nelectrochemically active surface area (ECSA) to isolate intrinsic catalytic activity.\n5. DFT calculations for amorphous regions: Incorporated additional simulations assessing the\ninfluence of amorphous domains on C–N coupling energetics.\n6. Post-reaction TEM: Added high-resolution TEM characterization of post-reaction catalysts\nto evaluate morphological stability.\nWe also acknowledge the Reviewer’s constructive request for further clarifications. In the following,\nwe provide detailed, point-by-point responses to each of the questions raised, and have revised the\nmanuscript accordingly to improve its clarity, rigor, and completeness.\nunclear whether such synergy is superior to a physical Cu + Co composite. Additional sample\nsynthesized from physically mixed Cu/Co electrodes or sequential Cu-on-Co deposition should be\ndiscussed.\nReply: We thank the Reviewer for this insightful comment regarding whether the observed tandem\nmechanism is truly a result of atomic-level mixing via co-sputtering, or if similar performance could\nbe achieved by a simple physical combination of Cu and Co.\nIn the revised manuscript, we have clarified this distinction. The co-sputtering process yields a Cu–Co\ncatalyst with atomic-level mixing, as confirmed by STEM–EDS mapping and homogeneous XAS\nsignatures, which ensures intimate proximity between Cu and Co sites. This structural feature is\nessential for enabling the cooperative interaction between *NH (Cu-favored) and *COOH/*CO (Co-\n2\nfavored) intermediates, as further supported by our DFT calculations showing reduced energy barriers\nat the Cu–Co interface compared to pure Cu or Co surfaces.\nTo probe this point, we conducted additional experiments using a physically mixed Cu/Co catalyst\nprepared by drop-casting (Fig. R4, depicted as Fig. S33 in the revised Supporting Information). Unlike\nthe co-sputtered catalyst, the drop-cast Cu/Co sample produced no C–N products; instead, only\nammonium product, formate, H , CO, acetate, ethanol, and propanol were detected. These results\n2\nindicate that Cu and Co act independently in the absence of atomic-level coupling, and do not function\nin a tandem relay mechanism when simply combined physically.\nPrior studies similarly underscore the importance of intimately coupled dual sites. For example, Cu–\nCo alloy nanoparticles encapsulated in N-doped carbon (CuCo@N-CNTs) were shown to balance CO\n2\nand NO ⁻ activation, delivering urea yields through cooperative C–N coupling.144 Likewise, sputtered\n3\nbimetallic catalysts have been demonstrated to create effective dual-site environments that facilitate\nelectrochemical C–N bond formation.16 These precedents reinforce our interpretation that the synergy\nobserved in our co-sputtered Cu–Co catalysts arises from atomic-level mixing, rather than from a\nsimple physical combination of metals.\nWe have added this discussion in the revised manuscript (Page 12, Lines 244-247). Together with our\nnew drop-cast Cu/Co control experiment, these results directly demonstrate that simple physical\nmixing of Cu and Co does not lead to C–N coupling, whereas co-sputtering provides atomic-level\nmixing that is essential for tandem synergy.\n“Importantly, this effect cannot be replicated by simple physical mixing: a drop-cast Cu/Co composite\nproduced only ammonium, formate, H , CO, and C alcohols (Fig. S33), with no detectable urea.\n2 2+\nThis finding demonstrates that atomic-level proximity achieved by co-sputtering is critical for\ncooperative interactions and the relay-type mechanism required for urea synthesis.”\nunderlying feature. The peaks should be re-fitted with appropriate constraints to ensure accurate\ndeconvolution.\nReply: We appreciate the reviewer’s observation regarding the Co 2p region. The broad main line and\nthe small underlying features are consistent with the well-known multiplet splitting and shake-up\nstructure of first-row transition-metal oxides/hydroxides, which can complicate deconvolution.\nBy comparing the Co 2p main line and satellite envelope of our catalysts with reference spectra in\n3/2\nliterature,68 we found that the Co 2p region is best described by a mixed Co(0) and Co(OH) signature.\nexplicitly including the appropriate multiplet and satellite components. The updated fits are shown in\nFig. R5 (adapted as Fig. 2 in the revised manuscript) and Fig. R24 (adapted as Fig. S12 in the revised\nSupporting Information).\n(n) (1-n) 1 2\n(C-D) Cu Co , and (E-F) Cu Co catalyst. The peaks of the high-resolution Co 2p spectra for all catalysts are deconvoluted\n1 1 2 1\nfor Co(0) and Co(II), where the Co 2p main line and satellite features align with reference spectra for Co(0) and\n3/2\nCo(OH) .68 These results shows that Co metal on the surface of catalysts are oxidized to Co(II) due to air exposure.71,72\n12 h reaction in the CO -saturated 0.2 M KHCO + 0.02 M KNO .\n2 3 2\nrationale is only superficially discussed, which should be strengthened.\nReply: We thank the Reviewer for this helpful suggestion. We agree that the mechanistic rationale for\nthe pH dependence of urea formation should be elaborated. In the revised manuscript, we have\nsubstantially strengthened this discussion by integrating experimental evidence with mechanistic\ninsights. At neutral pH, the balance of precursor availability, proton activity, and competing HER is\noptimized to favor C–N coupling:\n1. Precursor balance (*CO/*COOH & *NH ): Neutral bicarbonate media preserve dissolved\n2\nactivity to promote *CO/*COOH and *NH generation, unlike acidic (proton-dominates →\n2\n2. Local buffering. Neutral KHCO efficiently buffers cathodic alkalinization, stabilizing\n3\noxygenated intermediates such as *NCO and *NH CO without excessive *OH site blocking\n2\n(common in strongly alkaline electrolytes).142,143\n3. Catalyst stability: CV (Fig. R25, depicted as Fig. S34 in the revised Supporting Information)\nand ICP (Fig. R26, depicted as Fig. S35 in the revised manuscript) measurements show that\nneutral conditions suppress aggressive Cu/Co redox cycling (acidic) and irreversible\noxide/hydroxide formation (alkaline), preserving the Cu–Co interface.143,145\n4. Cu–Co interface synergy: The tandem mechanism requires the simultaneous availability of\n*CO/*COOH (most likely from Cu) and *NH (most likely from Co) (Fig. R4, depicted as\nthat the effective C–N coupling rate scales with their product. This is directly supported by our\nin situ SR-FTIR, which reveals co-existing *CO/*COOH and *NH signatures together with\n2\nurea-related carbonyl/CN bands only under neutral conditions, whereas monometallic controls\nlack one of the two precursor pools (Fig. R1, depicted as Fig. S42 in the revised Supporting\nInformation).\nTaken together, these factors explain why neutral pH uniquely favors urea formation: it maximizes\nprecursor coverages, suppresses HER, and maintains catalyst stability, thereby enabling efficient relay-\ntype C–N coupling at the Cu–Co interface. We have clarified this discussion and its connection to our\nin situ spectroscopic data in the revised manuscript (Page 12-13, Lines 264-308).\n“Beyond tuning the Cu:Co ratio, we also explored the effects of the ion-exchange membrane and pH\nof the catholyte and anolyte on the effectiveness of the C‒N coupling for the urea generation. Control\nexperiments were performed by adjusting the pH of both catholyte and anolyte to acidic, neutral, and\nalkaline conditions under continuous CO -gas flow (Table S6). No urea was detected after co-\n2\nreduction of CO and NO ⁻ at −1.2 V vs. RHE under either acidic (pH 3 & 5, K SO , Nafion membrane)\n2 2 2 4\nor alkaline conditions (pH 10 & 13, KOH, anion exchange membrane (AEM)). Under alkaline\nconditions, formate was the dominant product, acidic electrolytes suppressed C‒N coupling and\n60\nsteered the reaction toward H and ammonium generation (Fig. S36). Urea formation was consistently\n2\nobserved only when KHCO was used as buffer in both catholyte and anolyte, together with Nafion as\n3\nmembrane (Fig. 4F & Fig. S31). Notably, replacing Nafion membrane with AEM under identical\nconditions eliminated urea production (Table S6). These results indicate that controlled proton\navailability is critical for regulating surface protonation steps, consistent with a PCET pathway rather\nthan mediation by surface-bound hydrogen (*H).101,133,134 This conclusion is further supported by in\nsitu SR-FTIR measurements (Fig. S41), which show the absence of broad vibrational signatures\ncharacteristic of *H-related species in the 3300–3400 cm⁻1 region.135,137\nTo rationalize these trends, we consider the electrolyte-dependent mechanistic factors. Neutral\nKHCO maintains relatively high dissolved CO concentrations, in contrast to alkaline conditions\n3 2\nwhere CO is converted to HCO ⁻/CO 2⁻, thereby enabling robust *CO/*COOH formation.138 In\n2 3 3\nacidic electrolytes, *CO can form but HER dominates surface sites and mass transport, strongly\ninhibiting C–N coupling.139 For NO ⁻ reduction, neutral pH provides sufficient proton activity to drive\n2\n*NO hydrogenation to *NH without excessive HER, whereas limited proton availability in alkaline\n2\nelectrolytes slows hydrogenation steps and reduces *NH coverage.123,140,141 Furthermore, neutral\n3 2\nwithout severe *OH blocking that occurs in strongly alkaline media.110,142,143\nAnolyte variation further highlighted the importance of proton flux (Table S6). Maintaining the\ncatholyte at neutral KHCO but lowering the anolyte to pH 5 enabled detectable urea (FE 2 ± 0.2%,\n3\nyield rate 14 ± 2 mmol h‒1 g ‒1; Fig. 4F & Fig. S37), whereas excessive proton concentration (pH 3)\ncat\ndiminished CO availability and enhanced competing H /NH + pathways. This delicate balance in H⁺\n2 2 4\nconcentration highlights the nuanced interplay between CO activation and NO ⁻ hydrogenation, in\n2 2\nwhich optimal PCET-driven protonation enables co-coverage of *CO/*COOH and *NH .23,133,136,144\nvibrational bands under neutral conditions (Fig. S41-S42). Collectively, these results emphasize that\nurea formation requires not only an optimal Cu–Co surface but also finely tuned proton management\nacross both membrane and electrolyte compartments.\nThe role of electrolyte pH on catalyst stability is further reflected in CVs recorded in the non-catalytic\npotential window (Fig. S34). Acidic electrolytes promoted rapid Cu/Co redox cycling, consistent with\naccelerated dissolution.145,146 Whereas alkaline electrolytes induced broad, hysteretic oxidation waves\ncharacteristic of irreversible oxide/hydroxide formation.143 By contrast, neutral KHCO displayed\n3\nonly weak and broadened transitions, indicating moderate surface dynamics. These CV signatures\ndemonstrate that acidic and alkaline conditions drive aggressive redox activity and oxide formation,\ncompromising structural integrity, whereas neutral conditions preserve a stable Cu–Co interface.\nThe instability was corroborated by ICP analysis of Cu and Co dissolution under strongly acidic and\nalkaline environments (Fig. S35). In acidic media (pH 3), Cu and Co dissolution achieved 8 ± 2 wt.%\nand 17 ± 2 wt.%, respectively. In alkaline media (pH 13), dissolution was lower (3 ± 1 wt.% for Cu\nand 8 ± 2 wt.% for Co). Such losses significantly undermine the long-term stability and performance\n61\nof the catalysts. Finally, we observed that urea is prone to decomposition into ammonium and\nformate/formic acid under highly acidic (pH 3, K SO ) and alkaline (pH 13, KOH) conditions (Fig.\n2 4\nS38), especially at low concentrations.”\nFigure R25 | The percentage loss of Cu Co catalyst following CO and NO ⁻ electroreduction. (A) Percentage\n(n) (1-n) 2 2\nloss of Cu and Co metals using different catalysts. (B) Percentage loss of Cu and Co metals following CO and NO ⁻\n2 2\nelectroreduction using Cu Co catalyst in alkaline and acidic conditions. The concentrations of Cu and Co in the electrolyte\n1 1\nwere quantified using ICP-MS. The percentage loss was measured after 12 h reaction -1.2 V vs. RHE for a catalyst with\nan area of 0.5 cm2. Experiments in panel (A) were carried out in the CO -saturated 0.2 M KHCO + 0.02 M KNO solution\n2 3 2\nat -1.2 V vs. RHE. The error represents the standard deviation for at least three independent measurements. Experiments\nin panel (B) were conducted using 0.2 M K SO and 0.1 M KOH for acidic and alkaline anolytes, respectively. The pH of\n2 4\nacidic anolyte was adjusted by adding H SO . The error represents the standard deviation for at least three independent\n2 4\nmeasurements.\n1 1 2 2\n(B) enlarged view of panel (A). (C) Full CV normalized to ECSA and (D) enlarged view of panel (C).\nThe mechanistic role of electrolyte pH is further clarified by CV performed in non-catalytic potential windows (C,D). At\nneutral pH, only weak and broad redox features are observed, indicating that Cu and Co oxidation states are only mildly\nperturbed. This behavior suggests that bicarbonate buffering stabilizes surface oxides/hydroxides, suppressing extensive\nredox cycling and thereby maintaining catalyst integrity. In contrast, acidic electrolytes yield sharp, symmetric redox peaks,\nreflecting rapid Cu2+/Cu+ and Co3+/Co2+ transitions.145 Such pronounced redox cycling is consistent with accelerated Cu/Co\ndissolution observed by ICP analysis (Fig. S35) and explains the structural instability of the catalyst under strongly protonic\nenvironments.146 Under alkaline conditions, broad and intense oxidation waves with large hysteresis are evident,\ncorresponding to the formation of inactive Co and Cu oxides/hydroxides (e.g., Co(OH) , Co O , CuO, Cu(OH) ).143 The\n2 3 4 2\nsluggish reversibility reflects lattice restructuring, which destabilizes the catalyst surface. These electrochemical signatures\nhighlight that neutral KHCO uniquely moderates redox dynamics, preserving a mixed-valence Cu–Co surface optimal for\n3\ntandem C–N coupling.\nExperiments at pH 7 were performed in the CO -saturated 0.2 M KHCO + 0.02 M KNO solution. Acidic and alkaline\n2 3 2\nexperiments were carried out in CO -saurated 0.2 M K SO (pH adjusted with H SO ) and 0.1 M KOH for, respectively.\n2 2 4 2 4\narea (ECSA) to isolate intrinsic catalytic activity and decouple surface area effects.\nReply: We thank the Reviewer for this valuable suggestion. As recommended, we have normalized\nboth current densities and urea yield rates to the electrochemically active surface area (ECSA) to better\nreflect intrinsic catalytic activity and decouple surface area effects. These results are included in the\nFig. R22 (depicted as Fig. S22 in the revised Supporting Information) and Fig. R27-R28 (depicted as\nFig. S29 & S31 in the revised Supporting Information), in which urea yield rates are reported both per\ngram of catalyst and per cm ⁻2, and current densities (j) are expressed as mA cm ⁻2.\nECSA ECSA\nFigure R27 | i-t and i-V plots for CO RR and NO ⁻RR (normalized to ECSA). The i – t graphs of (A) Cu Co ,\n2 2 1 1\n(B) Cu, and (C) Co catalysts during electroreduction of CO and NO ⁻ at -1.2 V vs. RHE. Experiments were carried out in CO -\n2 2 2\nsaturated 0.2 M KHCO + 0.02 M KNO solutions.\n3 2\n2 2\ngraphs of Cu Co catalysts during electroreduction of CO and NO ⁻. (C-D) The i – t graphs of (C) Cu and (D) Co catalysts\n1 1 2 2\nduring electroreduction of CO and NO ⁻. (E) Stability test for Cu Co catalyst for 48 h. Panel (A, C-E) were conducted at\n2 2 1 1\n-1.2 V vs. RHE. All experiments were carried out in CO -saturated 0.2 M KHCO + 0.02 M KNO solutions.\n2 3 2\nam wondering whether the authors consider the defects or amorphous regions formed during\nsputtering, which may influence adsorption energetics.\nReply: We thank the Reviewer for raising this important point regarding the choice of surface models\nin our DFT calculations. We selected Cu(111) and Co(111) facets as model surfaces because they\nrepresent the most stable low-index planes and are widely used in computational studies of CO RR\n2\nand NO ⁻RR. This approach enables direct comparison with prior literature and highlights the intrinsic\n2\nactivity of Cu and Co toward *COOH, *CO, and *NH formation.\nenvironments that may influence adsorption energetics and catalytic behavior. In practice, however,\nsputtered catalysts typically form thin, relatively smooth films with fewer extended defects than oxide-\nderived systems.45 In such cases, enhancement of C–N coupling is often achieved by introducing a\nsecondary component to form binary alloys, which tune adsorption energetics and selectivity.16\nTo address the Reviewer’s concern, we attempted to construct amorphous surface models for DFT.\nDespite several trials, these configurations could not be stabilized or converged. As an alternative, we\nanalyzed previously developed hybrid Cu–Co models with different surface orientations (Cu(111),\nCo(111), and a Cu/Co interface). These models exhibited local structural disorder resembling\namorphous features, and their calculated energy profiles revealed distinct variations in the C–N\ncoupling pathway, reinforcing the tandem role of Cu and Co sites.\nWhile incorporating these effects is beyond the scope of the present study, we have acknowledged this\nlimitation in the revised manuscript (Page 17, Lines 418-421). Specifically, our current models may\nunderestimate the contributions of defect/amorphous sites, which could further lower activation\nbarriers and stabilize key intermediates. Future work employing large-scale simulations or machine-\nlearning-assisted approaches will be better suited to capture the structural complexity of sputtered\ncatalysts.\nFinally, we emphasize that our experimental evidence (XAS, XPS, Raman, IR) consistently indicates\nmetallic Cu and Co as the dominant active phases under reaction conditions. Thus, while surface\ndisorder is expected, the (111)-plane and Cu–Co interface models provide a reasonable baseline for\nmechanistic interpretation, particularly in correlating Cu–Co interfacial synergy with the observed\nactivity and selectivity. The updated discussion can be seen in the revised manuscript (Page 17, Lines\n418-421):\n“We acknowledge that our present models likely underestimate the contribution of defect-rich or\namorphous sites, which could further lower activation barriers and stabilize key intermediates.148\nFuture studies employing large-scale simulations or machine-learning-assisted approaches will be\nneeded to better capture the structural complexity of sputtered catalysts.149”\nmissing.\nReply: We appreciate the reviewer’s suggestion. We have added post-reaction HRTEM for Cu₁Co₁,\nas well as pure Cu and pure Co controls (Fig. R29; shown as Fig. S7 in the revised Supporting\nInformation). These images confirm that the dominant lattice fringes are retained after electrolysis and\nreveal only minor surface oxides.\nRegarding XRD, reliable post-reaction patterns could not be obtained because the catalysts are thin\nfilms on carbon fiber paper (CFP) at low mass loading. The intense, broad CFP reflection and high\nbackground overwhelm the weak diffraction from the nanostructured coating; despite extended\nacquisitions, only the carbon peak was detected. As XRD is therefore non-informative for this sample\ngeometry, we rely on HRTEM/SAED, XPS, and XAS (XANES/EXAFS) to determine phase/oxidation\nstates before and after reaction.\nThe additional post-reaction HR-TEM characterization is provided in Fig. R29 and has been\nincorporated into the main manuscript (Page 5, Lines 76-78) as follows:\n“Importantly, these well-defined crystal facets are preserved even after electrochemical testing, as\nevidenced by the post-reaction TEM analysis (Fig. S7).”\n1 1\nCu Co , (B) Co, and (C) Cu catalysts. The dominant lattice fringes match those of the pre-reaction samples (metallic Cu\n1 1\nand Co), indicating that the bulk crystallographic structure is retained after electrolysis. 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MRS Communications (2025). https://doi.org:10.1557/s43579-025-00738-5\nStudy. physica status solidi (b) 258 (2021). https://doi.org:10.1002/pssb.202100185\n73"}, {"round": 2, "reviewer_comments": {"reviewer_1": "The authors have conducted substantial additional experiments to address the comments from the other reviewers and me. The data analysis and interpretation are now clearer, and the conclusions are better supported. I appreciate these efforts.\nBelow I have some further comments:\n1. In the Introduction part, the authors stated that “Cu suffers from limited urea selectivity and deactivation caused by strong adsorption of nitrogen intermediates. Cobalt, on the other hand, offers high conductivity, favorable *CO binding, and the ability to stabilize *H intermediates, thereby promoting downstream protonation or hydrogenation steps.” However, according to the experimental data presented in this study, it appears that Cu primarily stabilizes CO2RR intermediates (e.g. *CO/*COOH), whereas Co is in charge of stabilizing nitrogen-related intermediates (e.g. *NH2). This point should be clarified to avoid inconsistency.\n2. Fig. S19 indicates that a higher fraction of oxidized species is generated after electrochemical reduction compared to as-prepared catalysts. Although both the as-prepared catalysts and the post-reaction catalysts are characterized ex situ and catalyst oxidation in air is unavoidable, the reasons for such difference should be provided.\n3. The authors compared the CV features of different catalysts under different reaction conditions (Fig. S24). Please explain why the CV features corresponding to quasi-reversible catalyst redox peaks (not related to catalytic reactions) change under different electrolytes. Given that the same catalysts are used, why do peak shapes, positions, and areas vary?\n4. For the experiment of Raman spectroscopy, the use of 50% laser power is too high. It is recommended to lower the power to minimize the damage while still ensuring sufficient signal quality.\n5. In Fig. S36, the authors adjusted the pH of the anolyte to study the effect of proton availability. When the anolyte pH is changed, does the catholyte pH also change accordingly? Please clarify whether the performance difference could be attributed to this reason.", "reviewer_2": "The authors have addressed all my questions, and I have no further additional comments.", "reviewer_3": "The authors addressed all comments and therefore I am happy to recommend it for publication."}, "author_response": "Response to the first Reviewer’s comments\nComment: The authors have conducted substantial additional experiments to address the comments\nfrom the other reviewers and me. The data analysis and interpretation are now clearer, and the\nconclusions are better supported. I appreciate these efforts. Below I have some further comments:\nReply: We sincerely thank the Reviewer for the careful evaluation of our manuscript and the\nconstructive comments provided. We have carefully revised the manuscript to address these concerns\nas outlined below:\n Clarification of Catalyst Roles and Mechanism. The Introduction section has been revised\nto include working and alternative hypotheses based on prior studies. Our results support the\nalternative hypothesis that Cu predominantly stabilizes CO RR intermediates (*CO/*COOH),\n2\nwhile Co drives NO ⁻ reduction and stabilizes N-intermediates (*NH ), resolving the\n2 2\npreviously noted inconsistency.\n Surface Characterization. We provide an extended explanation for the higher fraction of\noxidized metal species in post-reaction samples (Fig. R1, depicted as Fig. S19 in the revised\nSupporting Information). This discussion now includes the roles of redox cycling, dynamic\nsurface restructuring, and susceptibility to re-oxidation upon ex situ analysis.\n Voltammetry Variability. We elaborate on why the quasi-reversible Cyclic Voltammetry\n(CV) features vary across different electrolytes (Fig. R2, depicted as Fig. S24 in the revised\nSupporting Information). The figure captions now include discussion on the influences of pH,\nspecific adsorption, double-layer structure, and ionic conductivity; we add concise notes in the\nfigure caption.\n Raman Power Optimization. We document our Raman power optimization, including control\nspectra at lower powers (10–25%) and rationale for not using laser power below 50% due to\ninsufficient signal-to-noise ratio, with a clarifying sentence in the revised Supporting\nInformation (Fig. R3, depicted as Fig. S21 in the revised Supporting Information).\n pH Control. We clarify the pH control experiment (Fig. R4, depicted as Fig. S36 in the revised\nSupporting Information) by reporting catholyte pH monitoring (Table R1, depicted as Table\nS6 in the revised Supporting Information) and attributing performance differences primarily to\nproton availability/transport across the membranes.\nWe trust that these revisions address the Reviewer’s concerns and significantly strengthen the\nmechanistic depth of the manuscript.\n2\n*CO/*COOH), whereas Co is in charge of stabilizing nitrogen-related intermediates (e.g. *NH ). This\n2\npoint should be clarified to avoid inconsistency.\nReply: We thank the reviewer for this helpful observation. Our original Introduction section\nsummarized a working hypothesis based on prior reports suggesting that Co can favor *CO/*H\nstabilization during individual CO reduction reaction, while Cu may suffer from strong N-\n2\nintermediate binding during pure NO ⁻ reduction reaction. However, after we did our experiments,\n3\nin situ characterization, and DFT calculations, indeed we found the opposite results, in which Cu improve\nthe CO adsorption and activation, while Co more into *NH stabilization.\n2 2\nTherefore, we revised the Introduction to present the prior statements as a working hypothesis,\nintroduce an alternative cross-assigned model, and preview our finding that Cu enhances CO\n2\nadsorption and *NH formation, while Co promotes NO ⁻ adsorption and stabilizes N-intermediates,\n2 2\nconsistent with our in situ/DFT results.\nThe revised manuscript (Page 3-4, Lines 30-59) presented as follows:\n“Copper has been widely studied due to its ability to facilitate CO and NO ⁻ activation, moderate\n2 x\nhydrogen binding, and tendency to form key intermediates such as *CO and *NH .1-6 However, Cu\n2\nsuffers from limited urea selectivity and deactivation caused by strong adsorption of nitrogen\nintermediates.7-11 Cobalt, on the other hand, offers high conductivity, favorable *CO binding, and the\nability to stabilize *H intermediates, thereby promoting downstream protonation or hydrogenation\nsteps.10,12-15\nThese complementary characteristics led us to a working hypothesis that combining Cu and Co could\ncouple their strengths—enhancing NO ⁻ reduction to reactive nitrogen species on Cu, while stabilizing\nx\n*CO and *H on Co—to accelerate C–N bond formation.10,16,17 Supporting this concept, recent work\non amorphous Cu–Co boride (a-Cu CoB ) demonstrated alleviation of strong N-based and *CO\n0.1 x\nintermediates binding on the catalyst surface and achieved a faradaic efficiency (FE) of ca 28% with\na urea yield rate of ca 310 μg h-1 mg -1 (5.2 mmol h-1 g -1) from CO and NO ⁻.18\ncat cat 2 3\nIn parallel, we also considered an alternative, cross-assigned tandem hypothesis—that Cu\npredominantly improves the CO adsorption and activation, stabilizing *CO/*COOH,19,20 whereas Co\n2\nprimarily promotes NO ⁻ activation and stabilizes N-intermediates.21,22 To address the ambiguity and\n2\nresolve the roles of the two metals, we explored tandem Cu‒Co catalysts prepared via co-sputtering.\nThis approach enables atomic-level mixing of two metals, providing different but complementary\nactive sites for key reaction steps. We anticipated that such atomic-scale integration would enhance\nPCET, promote the formation of key intermediates, thereby optimizing the C‒N coupling process.23-25\nelectrolyte conditions—including pH and nitrogen-source species—to evaluate urea formation and\nidentify optimal conditions. Guided by these results, we employed in situ spectroelectrochemical\ntechniques—in situ Raman, synchrotron radiation Fourier-transform infrared spectroscopy (SR-\nFTIR), and XAS—to monitor the formation and evolution of reaction intermediates directly on the\ncatalyst surface under operating conditions. Finally, density functional theory (DFT) calculations\nwere performed under the optimized reaction conditions to provide atomic-scale insights into the roles\nof Cu and Co.\nTogether, these results support the alternative cross-assigned mechanism, rationalizing the observed\nselectivity where Cu predominantly stabilizes CO -derived intermediates (*CO/*COOH), and Co\n2\ndrives *NO ⁻ reduction and stabilizing *NH for C‒N coupling. This integrated experimental–\n2 2\ntheoretical approach aims to unravel the C–N coupling mechanism and guide the rational design of\nefficient catalysts for sustainable urea electrosynthesis.”\nComment 2: Fig. S19 indicates that a higher fraction of oxidized species is generated after\nelectrochemical reduction compared to as-prepared catalysts. Although both the as-prepared catalysts\nand the post-reaction catalysts are characterized as ex situ and catalyst oxidation in air is unavoidable,\nthe reasons for such difference should be provided.\nReply: We appreciate the reviewer for this insightful comment. As noted, both as-prepared and post-\nreaction catalysts were characterized as ex situ, and some degree of oxidation upon air exposure is\nindeed unavoidable. However, the higher fraction of hydroxide/oxide species observed in the post-\nreaction catalysts can be rationalized as follows:\n1. Surface reconstruction and reactivity. During electrochemical reduction, dynamic redox\ncycling induces partial restructuring of the catalyst surface (Fig. S39 in the revised Supporting\nInformation). This reconstruction produces a higher surface area with more reactive metallic\nsites that readily re-oxidize once exposed to the electrolyte. In addition, the electrochemical\nCO reduction reaction (CO RR) generates local OH⁻ at the surface of cathode. These high\n2 2\nOH⁻ concentration can react with the metal catalyst to form metal hydroxides (i.e., Cu(OH) \n2\nand Co(OH) ).26,27 Moreover, residual alkaline electrolyte (HCO ⁻/CO 2⁻, K+, residual OH⁻)\n2 3 2\nretained in the porous layer favors rapid formation of metal hydroxides upon brief contact with\ndissolved O during rinsing and drying.28,29\n2\n2. Thin surface layers dominate spectral changes. The increased oxide/hydroxide signal likely\nreflects nanometer-scale surface layers on the reconstructed high-area substrate rather than\nfrom wholesale bulk oxidation. Such ultrathin layers contribute disproportionately to the\nXANES line shape even in transmission/fluorescence geometry.30\nWe have revised the caption of Fig. R1 (depicted as Fig. S19 in the revised Supporting Information)\nto include this discussion, clarifying that while air exposure contributes to oxidation, the observed\ndifference most likely arises from structural/electronic modifications induced by the electrochemical\nreduction process.\n(A-F) the Cu K-edge and (G-L) the Co K-edge for Cu Co catalysts. The fittings indicate that both pre- and post-\n(n) (1-n)\nreaction Cu Co catalysts contain high concentrations of Cu and Co metals. Additionally, the spectra reveal that the\n(n) (1-n)\nCu Co catalysts undergo oxidation, forming oxide and hydroxide species as a result of air and electrolyte exposure.31,32\n(n) (1-n)\nThe pre-reaction catalyst refers to the as-prepared catalyst, whereas the post-reaction catalyst refers to the catalyst after a\n12 h reaction in the CO -saturated 0.2 M KHCO + 0.02 M KNO .We attributed the elevated oxide/hydroxide fraction in\n2 3 2\npost-reaction samples to several factors: (i) cathodic surface reconstruction during electroreduction, which exposes highly\nreactive metallic sites prone to rapid re-oxidation upon subsequent electrolyte or air exposure,33,34 (ii) the generation of\nlocal alkalinity during CO reduction, where elevated OH⁻ concentrations at the cathode surface can react with metal\n2\ncenters to form hydroxides such as Cu(OH) and Co(OH) ,26,27 (iii) the presence of residual alkaline environments (e.g.,\n2 2\nHCO ⁻/CO 2⁻, K+, OH⁻) within the catalyst layer that promotes hydroxidation of freshly exposed surface,28,29 and (iv) the\n3 2\ndisproportionate contribution of ultrathin surface oxides/hydroxides to XANES and Raman spectral features, such that\neven nanometer-scale re-oxidized layers produce pronounced signatures in ex situ measurements, as routinely reported in\nex situ studies.30\nconditions (Fig. S24). Please explain why the CV features corresponding to quasi-reversible catalyst\nredox peaks (not related to catalytic reactions) change under different electrolytes. Given that the same\ncatalysts are used, why do peak shapes, positions, and areas vary?\nReply: We thank the reviewer for raising this important point. As noted, the quasi-reversible peaks in\nthe CVs correspond to surface redox transitions of Cu/Cu–O(H) and Co/Co–O(H) species in the Cu–\nCo catalyst, rather than catalytic current. Although the catalysts are the same, the quasi-reversible\nredox peaks observed in CV are sensitive to several factors:\n1. pH-dependent proton-coupled electron transfer (PCET). The redox couples of transition\nmetals (e.g., Co2⁺/Co3⁺, Cu⁺/Cu2⁺) are proton-coupled and therefore highly sensitive to pH\naccording to Nernst equation.35-37 Variations in bulk electrolyte pH or buffering capacity\ndirectly alter the equilibrium potential (E ). For example, switching from Ar-saturated KHCO\neq 3\n+ KNO (pH 8, NO ⁻RR) to CO -saturated KHCO + KNO (pH 7, electroreduction CO and\n2 2 2 3 2 2\nNO ⁻) lowers the interfacial pH, causing the reduction peaks to shift to less negative (more\n2\npositive) potentials, as can be seen in Fig. R2I & Fig. R2L (depicted as Fig. S24I & Fig. S24L\nin revised Supporting Information). Furthermore, both CO RR and NO ⁻RR consume H+ and\n2 2\nH O, and generate OH⁻ ions at the cathode surface. This creates a significant interfacial pH\n2\ngradient that further shifts the Cu/Co redox equilibria from their bulk-derived values, directly\naccounting for the systematic potential shifts observed.38\n2. Specific anion adsorption and competitive binding. Electrolyte anions (HCO ⁻, NO ⁻, OH⁻)\n3 2\nexhibit specific adsorption on the Cu and Co sites, modifying the electric double layer and\nredistributing charge within the bimetallic interface.39-41 Moreover, the anions compete with\nwater molecules or protons for surface sites, affecting the fraction of the surface area available\nfor the intrinsic metal redox process. This leads to electrolyte-dependent changes to the peak-\nto-peak separation (ΔE ), peak symmetry, and peak intensity.42,43\np\n3. Electrochemical–chemical (EC) coupling and reconstruction. The redox peaks reflect a\ndynamic, potential-dependent surface phase transition that is coupled to the solution chemistry.\nFor instance, the presence of NO ⁻ (a mild oxidant) close to zero current densities can\n2\nchemically re-oxidize and thus stabilize the higher oxidation state (Co3+ or Cu2+), affecting the\nstarting surface state for the Cyclic Voltammetry (CV) scan.44,45 Moreover, the rapid local OH⁻\ngeneration facilitates the formation and dissolution of Cu(OH) /Co(OH) phases. These\n2 2\nprocesses render the waves quasi-reversible and cause peak broadening or area changes\ndepending on the electrolyte.26\n4. Tandem stabilization and interfacial effects. The Cu–Co interface modifies the\ndissolution/redeposition behavior of Co and stabilizes Cu⁺/Co2⁺ hydroxide species.46,47\nHowever, the degree of stabilization depends on the electrolyte: in CO -saturated KHCO\n2 3\n(CO RR), Cu–Co tends to maintain more reduced Cu sites, while in Ar-saturated KHCO +\n2 3\nNO ⁻ (NO ⁻RR), Cu–Co tandem stabilizes and suppress Co dissolution (Fig. R2F & Fig. R2I,\n2 2\ndepicted as Fig. S24F & Fig. S24I in revised Supporting Information). This leads to the\nobserved variations in peak intensities and integrated charges.\ncatalysts by tuning pH, adsorbate binding, local chemical reactions, and interfacial stabilization. The\nvariations in quasi-reversible redox peaks are therefore expected, even when the same catalyst is used,\nand they reflect electrolyte-dependent surface chemistry rather than changes in the bulk catalyst\nidentity.\nThus, while the intrinsic redox nature of the catalyst remains the same, the electrochemical\nenvironment provided by different electrolytes significantly impacts how these redox features are\nmanifested in CV. We have revised the Supporting Information to clarify this point and included\nadditional explanation in the caption of Fig. R2 (depicted as Fig. S24 in the revised Supporting\nInformation).\n1 1 2 2\nelectroreduction of CO and NO ⁻. CV curves for (A-C) HER, (D-F) CO RR, (G-I) NO ⁻RR, and (J-L) electroreduction\n2 2 2 2\nof CO and NO ⁻. CV measurements show that pure Cu exhibits reversible redox behavior (two oxidation and two reduction\n2 2\npeaks), whereas pure Co displays irreversible features (only one oxidation peak), attributed to Co2+/Co3+ dissolution into\nthe electrolyte.46,47 In contrast, the Cu–Co tandem stabilizes Co, suppresses dissolution, and modifies its electronic\nstructure, yielding a quasi-reversible system with enhanced stability and reusability in electrochemical reactions.48-50The\nquasi-reversible surface redox peaks vary with electrolyte due to: (i) pH-dependent proton-coupled electron transfer\n(PCET) equilibria in Cu/Cu–O(H) and Co/Co–O(H) couples (~59 mV per pH unit),35-37 (ii) specific adsorption and\ncomplexation of anions (HCO ⁻, NO ⁻, OH⁻), which modify the electric double layer structure and redistribute charge\n3 2\nwithin the Cu‒Co bimetallic interface,39-41 (iii) electrochemical-chemical (EC) coupling, where OH⁻ generated during\nCO RR/NO ⁻RR induces rapid hydroxidation or dissolution-redeposition, altering peak symmetry and charge, (iv)\n2 2\nchemical re-oxidation by NO ⁻ (a mild oxidant) near open circuit, which stabilizes the higher oxidation state (Co3+ or Cu2+)\n2\nand influences the initial surface state for the CV scan,44,45 and (v) interfacial affects in Cu–Co that suppress Co dissolution\nand stabilize Cu2⁺/Co2⁺ hydroxide species.46,47 Consequently, even with the same catalyst, electrolyte composition leads to\ndistinct changes in peak position, shape, and intensity across Ar-saturated KHCO , CO -saturated KHCO , Ar-saturated\n3 2 3\nKHCO + NO ⁻, and CO₂-saturated KHCO + NO ⁻.\n3 2 3 2\nAll experiments were conducted at scan rate of 20 mV s⁻1 and scan number = 5. The CV measurements for HER, CO RR,\n2 2 2 3 2\n0.2 M KHCO solution, Ar-saturated 0.2 M KHCO + 0.02 M KNO solution, and CO -saturated 0.2 M KHCO + 0.02 M\n3 3 2 2 3\nKNO solution, respectively.\n2\nrecommended to lower the power to minimize the damage while still ensuring sufficient signal quality.\nReply: We thank the reviewer for this constructive suggestion. We fully agree that minimizing laser\npower is important to avoid local heating or potential structural modifications. In our preliminary\nmeasurements, we attempted to acquire spectra of the Cu Co catalyst using 10% laser power;\n1 1\nhowever, the resulting spectra showed insufficient signal intensity and a very low signal-to-noise ratio,\nwhich hindered reliable peak assignment (Fig. R2D, depicted as Fig. S21D in the revised Supporting\nInformation). To balance spectral clarity with preservation of catalyst integrity, we optimized the laser\npower to 50%. Under these conditions, the spectra exhibited well-resolved and reproducible features\nwithout signs of laser-induced damage (e.g., no peak broadening, baseline drift, or irreversible changes\nupon repeated scans). We therefore selected 50% laser power as the optimal compromise between\nminimizing possible heating effects and ensuring sufficient signal quality.\nWe have revised Fig. R2D (depicted as Fig. S21D in the revised Supporting Information) accordingly\nto clearly document this optimization process and to explain why lower laser power settings could not\nbe used reliably.\nFigure R3 | Ex situ Raman characterization. Raman spectra of (A) pure Cu, (B) pure Co, and (C-D) Cu Co collected at\n1 1\nvarying laser power. (E-F) Comparison of as-prepared Cu Co catalysts acquired at 100% and 50% laser power. For\n(n) (1-n)\npure Cu, at 50% laser power, Cu O bands are observed at ~147, 219, and 634 cm⁻1.10,51 Upon increasing laser power,\n2\nadditional copper oxides/hydroxide features emerge. A similar trend is observed for pure Co; at 50% laser power, Co(OH)\n2\nbands are discernible at ~467, 516, 602, and 666 cm⁻1.52-55 However, with higher power, the spectrum evolves to Co O ,\n3 4\nshowing modes at ~189, 469, 510, 610, and 675 cm⁻1.52-55 The emergences of oxides/hydroxides in higher laser power may\nreflect improved signal-to-noise and/or laser-induced oxidation during Raman acquisition.52,56,57 For the as-prepared\nCu Co catalyst, spectra acquired at 50% laser power are dominated by Cu O and Co(OH) , whereas higher laser power\n(n) (1-n) 2 2\nreveals CuO bands at ~276 and 320 cm⁻1.51 At 100% laser power, Cu Co exhibits a mixture of Co O and Co(OH) features\n1 1 3 4 2\nat ~189, 470, 512, 600, and 664 cm⁻1,52-55 together with CuO bands at ~276 and 320 cm⁻1.51 Lowering laser power below\n50%, the metal peaks become barely detectable, likely due to insufficient signal intensity and a very low signal-to-noise\nratio.58 (Peak positions may vary by ±5–10 cm⁻1 due to local heating and instrumental parameters).\navailability. When the anolyte pH is changed, does the catholyte pH also change accordingly?\nPlease clarify whether the performance difference could be attributed to this reason.\nReply: We thank the reviewer for raising this important point regarding potential coupling between\nthe anolyte and catholyte pH.\nIn our two-compartment setup, the anode and cathode compartments were separated by an ion\nexchange membrane: an anion exchange membrane (AEM) for pH > 8 (alkaline electrolyte) or a\nNafion membrane for pH ≤ 8 (acidic/neutral electrolyte). These membranes function to isolate the bulk\nelectrolytes, ensuring the catholyte bulk pH remains distinct and stable regardless of the anolyte\nadjustments. We monitored the catholyte pH before and after the tests and observed only negligible\nvariations (within ± 0.2 units), indicating that the catholyte environment remained essentially stable\n(Table R1, depicted as Table S6 in the revised Supporting Information).\nThus, the observed performance differences observed in Fig. R4 (depicted as Fig. S36 in the revised\nSupporting Information) can be attributed primarily to changes in proton availability and transport\nacross the membrane (governed by the anolyte pH), rather than to significant shifts in catholyte bulk\npH. We have clarified this point in the revised Supporting Information to avoid misunderstanding, as\nwell as the main text (Page 13, Lines 293-295):\n“Crucially, while the anolyte pH was varied to modulate charge transport, the catholyte bulk pH was\nmaintained as intended, confirmed by negligible shifts (≤ ± 0.2 pH units) before and after testing.”\nTable R1 | The list of control experiments by changing catholyte and anolyte conditions.Control experiments were\nconducted to determine the conditions under which urea can be detected. All experiments were carried out using CO and\n2\n0.02 M KNO ⁻ as carbon and nitrogen sources at -1.2 V vs. RHE. The KHCO and K SO concentrations were kept at 0.2\n2 3 2 4\nM. The KOH concentration varied depending on the pH of solutions.\nNo Catholyte pH pH pH Anolyte pH pH Membrane Urea Other C‒N\nC C C A A\nSolution (C) after after Solution after Products\nCO reaction (A) reaction\n2\nflow\n3 3\n3 3\n2 4 2 4\n2 4 2 4 2 4\n2 4 2 4 2 4\n3\n2 4 2 4 3\n3 2 4\n3 2 4\n3 2 4\n3\n3\n1 1 2\nNO ⁻ electroreduction.The overall FE at CO -saturated catholyte: (A) alkaline (KOH, pH 10 & 13) and acidic (0.2 M\n2 2\nK SO , pH 3 & 5) conditions. Anolyte solution is the same base solution as catholyte, but without NO ⁻ present.\n2 4 2\nComparison of overall FEs using CO -saturated catholyte: (B) 0.2 M K SO (pH 5) + 0.02 M KNO and (C) 0.1 M KOH\n2 2 4 2\n(pH 13) + 0.02 M KNO at different anolyte conditions. The anolytes for panel (B) are 0.2 M K SO (pH 5) and 0.2 M\n2 2 4\nKHCO , while for panel (C) are 0.1 M KOH (pH 13) and 0.2 M KHCO . The acidic electrolyte was adjusted by the addition\n3 3\nof H SO . No C‒N products are detected in both alkaline and acidic catholytes. The alkaline catholyte suppresses HER\n2 4\ngeneration and mainly generates ammonium and formate.59 The acidic catholyte produces higher FE of H . Under both\n2\nconditions, the C product, ethanol, is consistently observed. However, the C product, n-propanol, is exclusively detected\n2 3\nunder acidic conditions. This underscore that proton (H+) availability is essential in facilitating the stabilization of reaction\nintermediates and subsequent C‒C coupling for C products.59-62 All experiments were conducted at -1.2 V vs. RHE.\n2+\nNafion membrane was utilized for conditions with pH ≤ 8, while anion exchange membrane (AEM) for alkaline conditions.\nThe error bars represent the standard deviation for at least three independent measurements.Note: Although the anolyte pH\nwas adjusted to regulate charge transport, the bulk catholyte pH remained stable, as evidenced by minimal changes (≤ ±\n0.2 units) before and after the measurements.\nH = hydrogen, CO = carbon monoxide, CH = methane, NH ⁺ = ammonium, HCOO⁻ = formate, CH COO⁻ = acetate,\n2 4 4 3\nCH OH = methanol, C H OH = ethanol, and C H OH = propanol.\n3 2 5 3 7\nthe activity of H2O. 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