Stable Cu^δ+ sites derived from analogous rectifying interface in AgCu biphasic aerogels for efficient urea electrosynthesis at low potential

Domain: physics
Published: 2026-01-02
Editorial outcome: major_revision

Abstract
Urea electrosynthesis through CO2 and NO3- coupling presents a promising alternative to energy-intensive industrial processes. However, intricate catalytic mechanisms and competitive reactions impede achieving high-efficiency C-N coupling. Herein, we constructed analogous rectifying interfaces in AgCu biphasic aerogels, by leveraging electronegativity differences to derive stable Cuδ+ active sites, which simultaneously promoted *NOH adsorption and *CO coverage, thereby improving C-N coupling dynamics, confirmed by both operando technique and theoretical calculations. Specifically, notable urea yield (54.8 mmol h-1 g-1 cat.) with Faradaic efficiency (FE, 36.6%) at a low potential (-0.52 V vs. RHE) was achieved using Ag67Cu33 in a flow-cell. As a proof-of-concept demonstration for practicability, Ag67Cu33 exhibited bifunctional electrosynthesis for urea (FE: 56.07%, yield: 104.6 mmol h-1 g-1 cat.) and HCOOH (FE: over 90%) at 40 mA cm-2 in two-electrode system, with durability over 60 h. This work provides an indicative rationale to construct active sites for C-N coupling, facilitating the development of urea electrosynthesis.

Introduction
In recent years, global research on greenhouse gas (GHG) reduction strategies has intensified, with a focus on finding innovative solutions for sustainable development. Approximately 12% of anthropogenic GHG emissions are attributed to agricultural production, particularly the synthesis of nitrogen fertilizers such as urea1–4. Traditionally, urea synthesis is dominated by the Bosch-Meiser process, which operates under extreme conditions (å180 °C, ~150 bar) and accounts for 2% of global energy consumption and 1.44% of CO2 emissions5,6. Therefore, it is crucial to explore sustainable methods for urea synthesis. With the rapid advancement of renewable energy technologies, electrocatalytic coupling of excess CO2 and harmful NO3- driven by renewable energy presents a promising alternative3,7,8. In 1995, Furuya et al. achieved a significant breakthrough by synthesizing urea for the first time using a Cu-supported gas diffusion electrode with CO2 and NO3- or NO2-9. Subsequently, Wang et al. successfully combined NO3- and CO2 into urea via a Vo-CeO2-750 catalyst, reaching a high yield of 943.6 mmol g-1 h-110. Recently, Liao et al. developed a PcNi-Fe-O catalyst and effectively obtained 0.164 g of high-purity urea in a scaled-up flow cell, achieving urea yield of 2.03 g h-1 g-1 cat.11. He et al. employed pulsed electrolysis on the synthesized Fe-TPP/CNTs catalyst, reaching a peak FE of 27.70%, rendering urea electrosynthesis more cost-effective and energy-efficient12. Despite considerable efforts, achieving optimal urea synthesis remains challenging due to complex catalytic mechanisms and multiple parallel competitive reactions (Supplementary Table 1)3, underscoring the need to develop high-efficiency electrocatalysts.

Copper (Cu)-based materials are widely recognized as highly effective catalysts for both CO2 reduction reaction (CO2RR) and NO3- reduction reaction (NO3RR) due to their suitable binding energies with C-intermediates and N-intermediates4,13,14. Consequently, Cu is regarded as one of the most promising materials for facilitating electrocatalytic urea synthesis5, whereas its poor selectivity in intermediate adsorption hinders C-N coupling efficiency13. To address this issue, substantial efforts have been made to improve the selectivity of Cu-based catalysts for specific products5,15,16. For example, by introducing amino substitutions, Wang et al. successfully optimized the electronic structure of CuPc-Amino catalyst, achieving great activity and stability in urea synthesis17. Therefore, modulating the electronic structures of Cu centers through techniques such as structuring interfaces18, incorporating dopants19, introducing defects20 has been suggested as a valuable strategy to effectively control the adsorption/desorption behavior of intermediates. Notably, previous studies have indicated that electron-deficient Cu plays a critical role in enhancing CO2RR and NO3RR by tuning the binding strength of intermediates21. For instance, Li et al. demonstrated that the Cu+-rich ultrathin 2D Cu2O nanosheets could effectively increase the adsorption capacity of *CO, facilitating the production of C2+ products15; Niu et al. designed a Cu2O/CuNRs/CF electrode that achieved rapid *NO3- adsorption, fast reduction of *NO3- to *NO, and highly selective conversion of *NO to *NOH22. Based on these findings, designing catalysts featuring Cuδ+ active sites has great potential to improve urea synthesis performance. However, Cuδ+ sites are easily reduced to Cu0 during electrolysis, resulting in catalyst deactivation, so there is an urgent need to develop strategies to stabilize Cuδ+ sites under negative potentials16. The rectifying interface effect in metal-semiconductor heterojunctions23 can cause charge redistribution and retain it near heterointerfaces by creating a built-in electric field driven by differences in work functions24, which provides a feasible approach for designing Cuδ+-based catalysts18.

However, fabricating catalysts with appropriate rectifying heterointerfaces requires careful consideration of many factors such as work functions, band structures, and dielectric constants, which are highly intricate and sensitive, presenting challenges in effectively controlling electron redistribution even through theoretical methods24. To address these challenges, we proposed using a second metal instead of the semiconductor to construct an analogous rectifying interface effect, which establishes a built-in electric field by utilizing electronegativity differences to induce electron transfer across interfaces. In addition, from the perspective of the material structure, metal aerogels, particularly bimetallic aerogels (e.g., BiSn25, PdAg26, etc.) emerge as promising candidates due to their three-dimensional porous structure, large surface areas, abundant active sites, and excellent structural durability27. Among these, Ag stands out as an ideal secondary metal due to its great electrical conductivity, thermodynamic immiscibility with Cu28,29 and its potential to induce electron-deficient Cu. Overall, the construction of AgCu biphasic aerogels with the analogous rectifying interface effect provides a feasible method to generate stable Cuδ+ sites and enhance electrocatalytic performance in urea formation.

Herein, we fabricated a series of AgxCu100-x biphasic aerogels featuring analogous rectifying interfaces via a one-step reduction using NaBH4. The analogous rectifying interface effect successfully generated stable Cuδ+ active sites near the interfaces, enhancing the critical C-N coupling dynamics, which resulted in an impressive urea yield of 54.8 mmol h-1 g-1 cat. with a high Faradaic efficiency (FE) of 36.6% using Ag67Cu33 catalyst at a low potential (-0.52 V vs. RHE) in a three-electrode flow cell, showing significant stability for 24 h, while also exhibiting great activity (yield: 22.59 mmol h-1 g-1 cat., FE: 32.4%) in an H-cell. Operando synchrotron radiation-Fourier transform infrared spectroscopy (SR-FTIR) and density functional theory (DFT) calculations revealed that the C-N coupling between *CO and *NOH constituted a pivotal step in urea electrosynthesis, while Cuδ+ sites effectively accelerated the reaction by simultaneously facilitating *NOH adsorption and improving *CO coverage. Specifically, compared to Cu0 sites, Cuδ+ sites exhibited a greater affinity for *NOH adsorption over *NHO, enabling successful C-N coupling with *CO to form *NOHCO through a thermodynamically spontaneous and kinetically accessible process, while *NHO was more prone to protonation, inhibiting the C-N coupling. Furthermore, Cuδ+ sites reduced the free energy barrier for CO2RR and stabilized *CO intermediate, which might increase *CO coverage on the surface for subsequent C-N coupling (Fig. 1). Additionally, to assess industrial feasibility, Ag67Cu33 was employed as a dual-function catalyst in a two-electrode flow cell that integrated cathodic urea synthesis with anodic formaldehyde oxidation reaction (FOR), which showcased great performance for both urea (FE: 56.07%, yield: 104.6 mmol h-1 g-1 cat.) and HCOOH production (FE: above 90%) at a high current density of 40 mA cm-2, maintaining satisfactory durability over 60 h. Our research highlighted the significant role of the analogous rectifying interface effect in AgCu biphasic aerogels and uncovered the crucial coupling reactions that enabled impressive urea synthesis, ultimately demonstrating the practical applicability of our catalyst in a two-electrode system.

Results
As depicted in Fig. 2a, a series of biphasic aerogel catalysts, AgxCu100-x (where x represented the theoretical mole fraction of Ag), were fabricated using a simple reduction method by introducing NaBH4 into a mixed solution of silver nitrate and copper nitrate (details were provided in the Methods section). The actual Ag and Cu content was determined by inductively coupled plasma mass spectrometry (ICP-MS), and the results (Supplementary Table 2) aligned with the expected stoichiometric ratios, indicating the complete reduction of metal ions. The X-ray diffraction (XRD) pattern of Ag67Cu33 (Fig. 2b) closely matched the standard reference patterns for pure Ag (PDF#04-0783) and Cu (PDF#04-0836), suggesting no oxide formation. No shifts in XRD peaks were observed in all catalysts (Supplementary Fig. 1), indicating that AgxCu100-x existed as distinct metallic Ag and Cu rather than forming an alloy30.

The morphology and structure of AgxCu100-x catalysts (Supplementary Fig. 2), as characterized by scanning electron microscopy (SEM), exhibited a typical aerogel with a three-dimensional porous network structure, which remained consistent across different molar ratios. Additionally, transmission electron microscopy (TEM) image of Ag67Cu33 (Fig. 2c) showed a network composed of interconnected nanoparticles, with each particle approximately 40 nm in diameter. High-resolution TEM (HR-TEM) image of Ag67Cu33 (Supplementary Fig. 3a) provided direct evidence for the separation of crystallized Ag and Cu along with the formation of Ag-Cu biphasic interfaces. The two distinct lattice spacings (Supplementary Fig. 3d-e), approximately 0.209 nm and 0.236 nm, corresponded to Cu(111) and Ag(111) planes, respectively, consistent with the selected area electron diffraction (SAED) pattern and XRD data. Furthermore, energy-dispersive X-ray (EDX) mapping (Supplementary Fig. 3b) revealed the phase-separated components in Ag67Cu33, with Ag and Cu concentrations of 61.62% and 38.38%, respectively (Supplementary Fig. 3c).

Subsequently, scanning transmission electron microscopy (STEM) combined with EDX spectroscopy, was utilized to investigate the Ag-Cu interfaces in more detail. The high-angle annular dark-field (HAADF-STEM) image (Fig. 2d), obtained by magnifying the white box in Fig. 2e, revealed a well-defined interface and atomic structures of metallic Cu and Ag with a crystallographic zone axis of [100], while the lattice spacings of 0.208 nm and 0.237 nm corresponded to cubic Cu(111) and cubic Ag(111), respectively (Supplementary Fig. 4). STEM-EDX mapping (Fig. 2e) and the line scan profile (Fig. 2f) further confirmed the distinct distribution of Ag and Cu, with a well-defined interface in biphasic aerogel, rather than the formation of an Ag-Cu alloy28, as expected based on the limited miscibility of Ag and Cu according to the binary phase diagram29. The surface elemental compositions and chemical states, analyzed by X-ray photoelectron spectroscopy (XPS) revealed the presence of Ag, Cu, O, and C (Supplementary Fig. 5). In the high-resolution spectra of the Ag 3d region (Fig. 2g), the peaks at 368.31 eV and 374.30 eV in Ag67Cu33 corresponded to Ag 3d5/2 and Ag 3d3/2 of Ag0, exhibiting a slight shift to lower binding energy compared to pure Ag (368.46 eV and 374.45 eV)26. Meanwhile, in the Cu 2p spectra (Fig. 2h), the Cu0/Cu+ peaks in Ag67Cu33, located at 952.46 eV (Cu 2p1/2) and 932.61 eV (Cu 2p3/2), shifted to higher binding energy compared to pure Cu (952.35 eV and 932.51 eV)14,31, indicating electron transfer from Cu to Ag and confirming the successful establishment of an analogous rectifying interface effect in AgCu biphasic aerogel. Although two peaks corresponding to Cu 2p1/2 and 2p3/2 of Cu2+ were present in both Ag67Cu33 (954.55 eV and 934.76 eV) and pure Cu (954.44 eV and 934.64 eV)32, the partial oxidation was attributed to surface exposure to air during ex-situ measurements30.

The urea synthesis efficiency of AgxCu100-x catalysts was evaluated in an H-cell with a three-electrode system (Fig. 3a) using controlled potential electrolysis, where the applied potentials ranged from -0.22 V to -0.62 V versus RHE (unless otherwise specified, all potentials were referenced to the reversible hydrogen electrode). Gas products (H2, CO) and liquid products (urea, NH3, NO2-) were detected using gas chromatograph (GC) and UV visible photometer (UV-Vis), respectively (see Methods section for details)6,14,18. Linear sweep voltammetry (LSV) curves for AgxCu100-x (Fig. 3b and Supplementary Fig. 6) showed that these catalysts exhibited activity for both CO2RR and NO3RR in CO2-saturated 0.5 M KHCO3 and Ar-saturated 0.5 M K2SO4 of 2000 ppm KNO3, respectively. Notably, AgxCu100-x demonstrated that higher current densities when CO2 and NO3- were introduced simultaneously, indicating a potential coupling reduction process8. Additionally, increasing the Cu content led to higher current densities, as shown in the i-t curves for AgxCu100-x (Supplementary Fig. 7), suggesting that Cu might be the active sites of the reactions. Figs. 3c, d illustrated the urea yield rate and FE of AgxCu100-x at various applied potentials (Supplementary Fig. 8), where Ag67Cu33 achieved the highest urea yield rate of 22.59 mmol h-1 g-1 cat. with a corresponding FE of 32.4% at -0.52 V. Notably, all AgCu catalysts exhibited better urea performance compared to pure Cu, confirming the critical role of the analogous rectifying interface effect in enhancing urea synthesis. Furthermore, the poor urea yield of pure Ag suggested that the improved performance was not due to the introduction of Ag alone. The FE of all possible products during the electrochemical urea synthesis on AgxCu100-x were measured (Supplementary Fig. 9 and Table 3), showing higher FE of NO2- and NH3 with increased Cu content or more negative potentials, which indicated that increased competition from NO3RR reduced urea FE33. In addition, AgCu catalysts exhibited lower FE for N-bearing by-products and H2 than pure Cu, especially Ag67Cu33, suggesting analogous rectifying interface effect might suppress competitive reactions.

Additionally, stability was another important parameter for evaluating catalyst performance. Recycling tests conducted on Ag67Cu33 at -0.52 V (Fig. 3e and Supplementary Fig. 10) revealed no significant decrease in current density, and the urea yield rate remained above 20.05 mmol h-1 g-1 cat. with an FE of over 27.8% after 6 continuous cycles. This great stability was corroborated by TEM, HRTEM, and XPS results after the stability experiments (Supplementary Fig. 11), which showed no obvious change throughout the electrolysis process. Moreover, post-test SEM (Supplementary Fig. 12) and XRD (Supplementary Fig. 13) analyses displayed no noticeable morphology and structural change in AgxCu100-x catalysts, except for pure Cu. The presence of Cu2O in Supplementary Fig. 13e likely resulted from air exposure after testing34, as the small size and high surface area of the pure Cu aerogel made it more susceptible to oxidation, while the introduction of Ag appeared to enhance the resistance to oxidation35. To further investigate the electrocatalytic activity of AgxCu100-x catalysts, the electrochemical active surface area (ECSA) was measured, as it played a crucial role in enhancing electrocatalytic performance36 (Supplementary Fig. 14). Despite the similar microstructures, Ag67Cu33 exhibited the largest ECSA, indicating a greater number of active sites available for catalytic reactions. Additionally, the turnover frequency (TOF) of AgxCu100-x, which signified the rate of electron consumption per unit time at an individual active site, was evaluated to gain deeper insights into their intrinsic catalytic activity37. As shown in Supplementary Fig. 15, the calculated n values for Ag, Ag80Cu20, Ag67Cu33, Ag50Cu50, and Cu were 4.09 × 10-10, 1.07 × 10-9, 1.21 × 10-9, 1.03 × 10-9 and 6.61 × 10-10, respectively, indicating that Ag67Cu33 possessed the highest number of active sites, which was consistent with the ECSA results. Moreover, Ag67Cu33 demonstrated the highest TOF value across the entire applied potential range, confirming its better intrinsic catalytic activity and thereby enhancing urea synthesis performance.

Isotope-labeling experiments were conducted in CO2-saturated 0.5 M KHCO3 electrolytes containing 2000 ppm K15NO3 or K14NO3 to verify the origin of the urea produced via the CO2 and NO3- coupling reaction. As shown in Fig. 3f, the 1H-NMR spectrum of the electrolyte with K15NO3 after electrochemical testing exhibited characteristic double peaks at 5.70 ppm and 5.55 ppm, consistent with the standard 15NH2CO15NH2. In contrast, a single peak at 5.63 ppm corresponded to the formation of 14NH2CO14NH2 within the electrolyte containing K14NO3, confirming that the urea produced in this work originated from the electrocatalytic coupling of NO3- and CO2. Moreover, the absence of urea in electrolytes lacking either CO2 or KNO3 (Fig. 3h and Supplementary Fig. 16) further supported that the urea synthesis was driven by the co-reduction of CO2 and NO3-, ruling out any potential impurities. Subsequently, the 1H-NMR method was also used to validate the accuracy of the urea yield determined by the diacetyl monoxime method. Based on the standard urea curve (Supplementary Fig. 26), the comparable urea yields from both methods confirmed the reliability of our results (Fig. 3g and Supplementary Fig. 17). Compared with previously reported catalysts (Supplementary Fig. 18 and Table 4), our work demonstrated great performance, with an impressive urea yield rate and high FE at low potentials.

To enhance the catalytic performance for electrochemical urea synthesis, we conducted electrochemical tests in a three-electrode flow cell (Fig. 4a and Supplementary Fig. 19a) under applied potentials ranging from -0.22 V to -0.62 V. Previous studies showed that flow cells significantly improved electrocatalytic performance in reactions such as NO3RR and CO2RR38. Due to the more competitive CO2RR in flow cells39, a higher concentration of NO3- was required to facilitate the C-N coupling reaction. As the concentration of KNO3 increased (Supplementary Fig. 19), Ag67Cu33 achieved a higher urea yield rate along with an increased FE, reaching its optimum at 0.5 M KNO3. However, excessively high KNO3 concentrations inhibited CO2 adsorption by promoting NO3- adsorption, thereby favoring competitive NO3RR and resulting in a decline in both urea yield rate and FE40. Thus, 0.5 M KNO3 was selected for subsequent electrochemical tests in the flow cell.

As shown in Fig. 4b, the significantly increased current density confirmed enhanced CO2RR activity. When CO2 and NO3- were introduced simultaneously, there was a notable increase in current density, although NO3RR became increasingly competitive at potentials beyond -0.62 V. Remarkably, Ag67Cu33 achieved a maximum urea yield rate of 54.8 mmol h-1 g-1 cat. with an FE of 36.6% at -0.52 V through potentiostatic tests (Fig. 4c and Supplementary Fig. 20), surpassing previously reported results (Fig. 4e and Supplementary Table 5). Additionally, as more negative potentials were applied, an increasing production of NO2- and NH3 was observed (Fig. 4d and Supplementary Fig. 21), consistent with the LSV curves. A long-term stability test (Fig. 4f and Supplementary Fig. 22) was conducted at -0.52 V to assess the intrinsic stability of Ag67Cu33. Notably, both the urea yield rate and FE remained above 50.1 mmol h-1 g-1 cat. and 36.9%, respectively, over a 24-hour period, while the SEM image and XRD pattern (Supplementary Fig. 23) of Ag67Cu33 after 24 h of electrocatalysis confirmed that the catalyst maintained its morphology and structural integrity. Furthermore, ICP-MS was utilized to analyze the dissolution conditions of the catalyst. As illustrated in Supplementary Fig. 24, trace amounts of Ag and Cu were detected in the post-electrolytes, with the final catalyst mass on carbon paper stabilizing at 0.97 mg (closely aligning with the calculated loss of 0.95 mg from the electrolytes). Notably, given the higher molar ratio (Ag:Cu = 71.9:28.1) of the post-test catalyst and the greater loss of Cu in the post-electrolytes, the slight catalyst loss could be attributed to surface oxidation at OCP when refreshing the electrolytes and subsequent dissolution under reductive potentials41. Despite a minor leaching during the stability tests, the catalyst remained largely intact after testing, confirming its satisfactory electrochemical and structural stability.

To gain mechanistic insights into the high activity of Ag67Cu33 in electrocatalytic urea synthesis, we employed synchrotron radiation X-ray absorption near edge structure (XANES) spectroscopy, operando Raman spectroscopy and synchrotron radiation-Fourier transform infrared spectroscopy (SR-FTIR). In regard to the XPS analysis results of Ag67Cu33 (Supplementary Fig. 11), it was evident that Ag existed in a zero-valent state, while Cu showed minor oxidation. Therefore, ex-situ synchrotron radiation XANES (Supplementary Fig. 31a) was employed to assess the chemical state of Cu in Ag67Cu33 following LSV activation and potentiostatic tests (from -0.22 V to -0.62 V). In the XANES spectrum of Cu L-edge at OCP (Supplementary Fig. 31b), the peaks at 931.9 eV and 951.8 eV were assigned to the 2p3/2 and 2p1/2 states of metallic Cu, while a minor peak indicative of Cu-O bonding was also observed33, likely arising from surface oxidation during the catalyst loading process. Notably, after electrochemical treatment, the spectral features exclusively exhibited metallic Cu peaks, without oxide signatures. Furthermore, operando Raman spectroscopy was utilized to further investigate this transformation. As shown in Fig. 5a, several vibrational modes observed at OCP were attributed to CuO (319 and 614 cm-1)16 and Cu2O (217 and 635 cm-1)22, resulting from surface oxidation following air exposure. During continuous LSV measurements to activate the catalyst, these peaks gradually diminished and disappeared after 10 min, indicating the complete reduction of CuO and Cu2O to metallic Cu. Furthermore, no peaks corresponding to either Ag or Cu oxides were detected during subsequent potentiostatic tests, suggesting Ag67Cu33 retained its metallic Cu and Ag composition throughout the electrochemical process (Supplementary Fig. 33), aligned well with the XANES results (Supplementary Fig. 31b).

Subsequently, operando SR-FTIR measurements (Supplementary Fig. 34) were conducted to monitor the evolution of reaction intermediates during the electrochemical process, with infrared signals collected between -0.22 V and -0.62 V using post-background subtraction for accuracy. At OCP, there was no significant interference from electrolytes or other species, which facilitated the detection of chemical signals. As illustrated in Fig. 5b, peaks observed at 1984 cm-1 and 2148 cm-1 corresponded to the stretching modes of *NO and *CO respectively42, indicating the successful absorption and co-activation of NO3- and CO2 on the catalyst. Additionally, the band at 1178 cm-1 was attributed to C-O stretching vibrations, suggesting CO2RR was occurring1,5. Moreover, a stretching vibration band associated with C-N emerged at 1452 cm-1 alongside -NH2 rocking vibrations at 1090 cm-17,43, peaking at -0.52 V, which was consistent with experimental results, providing crucial evidence for successful urea formation. It was noteworthy that an evident infrared band around 1686 cm-1 could be ascribed to *COOH and *CONH intermediates1,42,44, whose variation pattern with applied potentials agreed well with experimental observations, indicating that these intermediates, particularly *CONH, played a critical role in subsequent urea generation. Additionally, an infrared band associated with H2O vibration45 appeared within the range of 1640-1670 cm-1; thus control experiments using various electrolytes were performed via SR-FTIR to detect generated intermediate species while minimizing interference from H2O (Fig. 5c). Combined with electrocatalytic experiment results (Supplementary Fig. 16), when no CO2 gas was present, the infrared band at 1643 cm-1 should be assigned to H-O-H vibration; however, upon injecting CO2 gas, which led to the synthesis of either CO or urea, the infrared band broadened slightly while exhibiting a minor blue shift but retained a shoulder peak around 1643 cm-1, which further corroborated the band at 1686 cm-1 belonged to *COOH and *CONH intermediates38. Meanwhile, infrared vibrations associated with other products in CO2RR or NO3RR, including CO, NO2- and NH346–48, were also successfully tracked.

To gain a comprehensive understanding of the intrinsic mechanisms underlying urea synthesis and the role of analogous rectifying interface effect in AgCu biphasic aerogels for C-N coupling, density functional theory (DFT) calculations were performed (Supplementary Data 1). Given that Ag67Cu33 exhibited phase separation with well-defined interfaces and distinct Ag(111) and Cu(111) orientations, as indicated by the experimental results (Fig. 2), an Ag(111)-Cu(111) model with an interface was constructed (Supplementary Fig. 35). The Bader charge analysis of the Ag(111)-Cu(111) model was conducted to assess electron redistribution around the Ag-Cu interface (Fig. 5d). As predicted by the analogous rectifying interface effect, a built-in electric field arose between Ag and Cu due to their differences in electronegativity, resulting in the transfer of electrons from Cu to Ag, which led to the formation of stable electron-deficient Cuδ+ sites near the interface, consistent with the XPS results. Notably, Cuδ+ was identified as an active site for electrocatalytic urea synthesis, as well as for CO2RR19 and NO3RR20, based on previous studies and our computational results. The free-energy diagram for urea synthesis on Ag67Cu33 (Fig. 5e, Supplementary Fig. 36-39 and Table 6) showed that CO2 and NO3- could be adsorbed and activated with a low energy barrier. Importantly, compared to pure Cu0 (Supplementary Fig. 40-42 and Table 7), *NO tended to undergo protonation to form *NOH rather than *NHO on Cuδ+ sites, which played a key role in successful urea generation. Specifically, on Ag67Cu33, *NOH preferentially coupled with *CO to form the key intermediate *NOHCO via a thermodynamically spontaneous process (ΔG = -0.63 eV), and the transition state (TS) energy barrier (Supplementary Fig. 37) for the first C-N coupling (0.67 eV) was significantly lower compared to that required for the formation of *CO + *NHOH (0.95 eV), making C-N coupling more favorable from both thermodynamic and kinetic perspectives than the alternative route involving *NOH protonation. Following the second NO3- adsorption, the system successfully enabled the second C-N coupling and yielded the intermediate *HNCONO through an energetically downhill step (ΔG = -0.38 eV) with a relatively low TS energy barrier (0.87 eV), ultimately facilitating the efficient synthesis of urea. In contrast, on pure Cu surfaces, *NHO protonated to *NHOH more easily than undergoing C-N coupling due to the higher free energy barrier for the latter. Unfortunately, C-N coupling still failed to occur in subsequent reaction steps, resulting in poor urea production with CO and NH3 by-products instead (Supplementary Fig. 9q). To further investigate the impact of electron deficiency in Cu on intermediate adsorption, we calculated charge density distributions (Fig. 5f and Supplementary Fig. 46) and performed Bader analysis (Supplementary Table 10) for adsorbed species (*NO, *NHO, and *NOH) on Ag67Cu33 as well as pure Cu. Considering the electronegativity difference between H (2.2) and Cu (1.9), N and O were more likely to accept electrons from Cu, leading to the formation of the *NHO intermediate, where both N and O were in direct contact with Cu. However, it might be challenging for Cuδ+ sites to supply enough electrons, so another combination was adopted to form a more stable *NOH intermediate where H atom could transfer more electrons to O and N. Additionally, Cuδ+ sites not only reduced the energy barrier for CO2 conversion to *CO (ΔG = 0.44 eV), but also exhibited a higher endothermic barrier for *CO desorption (ΔG = 0.46 eV), enhancing CO activation and stabilizing *CO on the surface to improve its coverage36, which could promote further C-N coupling and facilitate urea synthesis, as indicated by Gibbs free energy changes in fundamental steps during CO2RR towards CO (Fig. 5g, Supplementary Fig. 47 and Table 9). Subsequently, the projected density of states (PDOS) of d orbitals further supported these conclusions (Fig. 5h), which exhibited that the d-band center of Cuδ+ sites shifted towards the Fermi level compared with pure Cu or pure Ag. According to d-band center theory49, an upshifted Ed of metal surface could enhance the binding strength between metal surface and adsorbed species, consistent with Gibbs free energy change results. Additionally, as shown in Supplementary Fig. 43-45 and Table 8, the free-energy diagram of urea synthesis on Ag(111) revealed that although *NHO readily coupled with *CO at a low energy barrier to form the first C-N bond (*CONHO), the high adsorption energy of CO2 and the weak binding strength of *CO (Fig. 5g) would result in insufficient *CO surface coverage. Furthermore, the limited active sites led to the poorest intrinsic catalytic activity (Supplementary Fig. 14-15), thereby collectively contributing to the restricted current densities and urea yield rates observed for Ag (Supplementary Fig. 7-8). To further verify the importance of the analogous rectifying interface effect, we physically blended pure Ag and pure Cu aerogel to create an Ag-Cu catalyst (Ag2Cu1) without interfaces (Supplementary Fig. 53). However, its urea synthesis performance (Supplementary Fig. 54) was significantly inferior to that of Ag67Cu33, with no notable enhancement observed, thereby emphasizing the essential importance of the analogous rectifying interface effect within Ag-Cu biphasic aerogels. Furthermore, by employing DFT calculations on NO3RR (Supplementary Fig. 48a, 49 and Table 11), HER (Supplementary Fig. 48b, 50 and Table 12), and CO2-to-C2+ (Supplementary Fig. 51-52 and Table 13), a comprehensive comparison was carried out among these potential competing reactions, which demonstrated that Cuδ+ active sites possessed a pronounced preference for C-N coupling, thereby promoting the formation of urea with satisfactory selectivity. Overall, the analogous rectifying interface effect (Supplementary Fig. 55) facilitated the formation of stable electron-deficient Cuδ+ active sites in AgCu biphasic aerogels, which might not only increase the affinity for *NOH and promote C-N coupling with *CO but also enhance the surface coverage of *CO through facile CO2 activation and *CO stabilization, ultimately accelerating C-N coupling dynamics and improving urea synthesis efficiency.

In order to further investigate the industrial potential of our catalyst, we designed a two-electrode flow cell that integrated urea synthesis with formaldehyde oxidation reaction (FOR), as Ag-Cu catalysts showed promising performance in FOR50. Thus, Ag67Cu33 was employed as a dual-function catalyst at both the cathode and anode in this system (Fig. 6a). This approach not only reduced the onset potential of the anode by replacing the sluggish oxygen evolution reaction (OER, å1.23 V vs. RHE) with a low thermodynamic potential process (HCHO + 2OH- → HCOO- + 1/2H2 + H2O + e-, E = -0.22 V vs. RHE)3,51, but also enabled the conversion of excess CO2 and waste NO3- into sustainable urea while producing high-value HCOOH. UV-vis spectroscopy and 1H NMR spectroscopy were employed to detect urea (Supplementary Fig. 25) and HCOOH (Supplementary Fig. 30), respectively. As shown in Fig. 6b, the LSV curves exhibited two distinct oxidation peaks in the absence of HCOH, corresponding to the Cu0/CuI and CuI/CuII oxidation processes50. Notably, when a high concentration of HCOH (0.6 M) was introduced into the anolyte, these oxidation peaks diminished due to the preferential oxidation of HCOH (-0.22 V RHE), occurring at lower potentials than Cu0/CuI and CuI/CuII transitions (approximately 0.4 and 0.7 V RHE, respectively). Furthermore, incorporating 0.6 M HCOH in the anode and 0.5 M KNO3 in the cathode resulted in a significant increase in current density, reaching 10 mA cm-2 at 0.34 V, a drop of approximately 1.5 V compared to OER||CO2RR, confirming the successful implementation of this two-electrode system.

After activization in an H-cell, Ag67Cu33 was directly employed as both the cathode and anode for subsequent constant current experiments (ranging from 15 to 45 mA cm-2). Specifically, as the current density increased (Supplementary Fig. 56), both the yield rate and FE of urea gradually rose (Fig. 6c), reaching maximum values at 40 mA cm-2 (104.6 mmol h-1 g-1 cat. and 56.07%, respectively). Concurrently, the yield rate of HCOOH continued to increase, achieving 3.09 mmol h-1 mg-1 cat. at 45 mA cm-2, although its FE experienced a slight decline, remaining above 90% (Fig. 6d). Additionally, SEM images and XRD results of both the anodic and cathodic catalysts after testing (Supplementary Fig. 57) demonstrated Ag67Cu33 maintains its aerogel morphology, along with metallic Ag and Cu phases, aligning well with the LSV curves. Further cycling tests were conducted at 40 mA cm-2 to assess the stability of Ag67Cu33 (Supplementary Fig. 58). The results indicated that both urea yield rate and HCOOH FE exhibited slight fluctuations, but consistently remained above 95.3 mmol h-1 g-1 cat. and 90.6%, respectively, over ten cycles (Fig. 6e). Although there was a minor increase in potential during each cycle, likely due to reactant depletion, the potential could be restored to its initial value by refreshing the electrolytes, demonstrating strong stability. Based on these findings, a current density of 40 mA cm-2 was selected to investigate the industrialization potential through a long-term stability test (Supplementary Fig. 59-60), with electrolyte refreshment every 10 h. As depicted in Fig. 6f, Ag67Cu33 exhibited well stability in both urea synthesis and FOR throughout the experiment. SEM images, XRD results and XPS spectra obtained after 60 hours of testing (Supplementary Fig. 61) were nearly identical to those before testing. Furthermore, the catalyst leaching at the cathode (maintaining 0.44 mg) was comparable to that observed in a three-electrode flow cell41, while the anode (retaining 0.43 mg) exhibited slightly higher Ag loss, which were attributed to the slight oxidation and dissolution processes occurring under anodic conditions50 (Supplementary Fig. 62). Owing to the intrinsic robust structure of metal aerogels and the enhanced oxidation resistance provided by Ag27, there was no rapid or extensive dissolution of the catalyst, thereby preserving its performance throughout the stability tests. All these findings further corroborated the great electrocatalytic activity and stability of our catalyst. In conclusion, the dual-functional Ag67Cu33 catalyst demonstrated significant electrocatalytic performance and stability for both urea synthesis and FOR within our two-electrode system at low cell voltages, significantly highlighting its potential for practical application.

Methods
Chemical and materials
Copper nitrate trihydrate (Ⅱ) (Cu(NO3)2·3H2O, AR), silver nitrate (Ⅰ) (AgNO3, 99.8%), sodium borohydride (NaBH4, 98%), potassium bicarbonate (KHCO3, 99.5%), potassium sulfate (K2SO4, 99%), potassium chloride (KCl, 99.8%), nitrate of potash (KNO3, 99%), Nafion solution (5 wt.%), urea (CO(NH2)2, 99%), iron chloride (Ⅲ) (FeCl3·6H2O, 99%), thiosemicarbazide (MTSC, 98.5%), diacetyl monooxime (DAMO, >98%), sulfuric acid (H2SO4, 95%-98%), phosphoric acid (H3PO4, >85 wt.% in H2O), potassium nitrate-15N (K15NO3, 99%), 15N-urea (CO(15NH2)2, 99%), dimethylsulfoxide-d6 (d6-DMSO, 99.9% D.), sodium nitrite (NaNO2, 99%), sulfonamide (C6H8N2O2S, 99%), N-(1-naphthyl)-ethylenediamine dihydrochloride (NETH, 98%), ammonium chloride (NH4Cl, 99.8%), potassium sodium tartrate tetrahydrate (C4H4KNaO6·4H2O, 99%), Nessler Reagent, ethyl alcohol (C2H5OH, 99.9%), nitric acid (HNO3, 65%-68%), potassium hydroxide (KOH, 90%), formaldehyde solution (36 wt.% in H2O), deuteroxide (D2O, 99.9% D.), dimethylsulfoxide (DMSO). All reagents were purchased from Sigma Aldrich Corporation and used without further purification. Carbon papers were purchased from Hesen company (Shanghai, China). High-purity deionized water (resistance 18 MΩ cm-1) were used to prepare aqueous solutions in all experiments.

Synthesis of the AgxCu100-x aerogels
The AgxCu100-x aerogels were prepared via a NaBH4-induced in situ reduction process. Taking the synthesis of Ag67Cu33 aerogel as an example, NaBH4 (50 mM, 25 mL) was homogeneously dissolved in deionized water in the ice-water bath, then the mixed aqueous solutions of Cu(NO3)2·3H2O (50 mM, 1.67 mL) and AgNO3 (50 mM, 3.33 mL) were rapidly dropped into the NaBH4 solution, and the color of the solution turned black instantly. After keeping stirring in the ice-water bath for 1 hour, Ag67Cu33 aerogel was collected by centrifugation and washed with water and ethanol for three times, and finally obtained by freeze-drying procedure. The AgxCu100-x aerogels were acquired using the aforementioned synthetic procedures by varying the molar ratio of Cu(NO3)2·3H2O and AgNO3 (Supplementary Table 2). The pure Ag and pure Cu aerogel catalysts synthesized using the identical method, with the only difference being the addition of either AgNO3 or Cu(NO3)2·3H2O.

The AgCu-mixture catalyst was synthesized by physically blending pure Ag and pure Cu aerogel (obtained from the aforementioned method) in a molar ratio of 2:1, ensuring a homogeneous distribution.

Materials Characterizations
The morphologies of the prepared electrocatalysts were characterized by cold field-emission scanning electron microscopy (FESEM, SU8220), transmission electron microscopy (TEM, JEOL, JEM-2010). The high-resolution TEM (HR-TEM) images and energy dispersive spectroscopy (EDS) mapping images were recorded using a Talox F200X (Thermo Fisher Scientific, America). Atomic-resolution HAADF-STEM was acquired using an aberration-corrected JEM-ARM300F operated at 300 kV. Powder X-ray diffractor (XRD, TTR-III), X-ray photoelectron spectrometer (XPS, Thermo Scientific ESCALAB 250Xi) with an X-ray source (Al Kα, hν=1486.6 eV, all the peaks were calibrated with the C 1s spectrum at a binding energy of 284.8 eV), synchrotron radiation X-ray absorption near edge structure (XANES) and Raman spectrometer (LabRam HR) with a 532 nm excitation laser were performed to study the structures and chemical compositions of electrocatalysts. Inductively coupled plasma atomic emission spectrometer (ICPAES, iCAP 7400) was conducted to determine the element ratios of the catalysts. UV visible photometer (UV-Vis, UV 3600), nuclear magnetic resonance (1H NMR, AVAVCE III HD 400 and JNM-ECZ600R/SI) and gas chromatograph (GC, GC 7900) equipped with a thermal conductivity detector and TDX-01 column were used for products quantification.

Electrochemical Characterization
The electrochemical measurements were conducted at 25 °C using the CHI660E electrochemical workstation, with continuous injection of CO2 into the electrolyzer or gas chamber during the electrocatalytic process. Data collection and processing were carried out using the Chi660e software. Furthermore, potentiostatic electrochemical impedance spectroscopy (EIS) was utilized to quantify the resistance of the electrochemical cell, ranging from 0.1 Hz to 200 kHz with an applied signal amplitude of 10 mVpp. The pH value of the catholyte was measured using a precision pH meter, revealing that its pH remained consistently at 7.4 ± 0.3 throughout the testing process. Prior to commencing the tests, we prepared the electrolytes (with the precise concentration detailed below) utilizing an electronic scale, beaker, graduated cylinder, and volumetric flask. The mixture was then thoroughly stirred to ensure complete dissolution, achieving a homogeneous solution. Subsequently, it was carefully transferred into a sealed container and stored at ambient room temperature (25 °C) for subsequent experimental use. And the measurements were only performed once. Reference electrode calibration required the immersion of the reference electrode in a buffer solution (pH 7), alongside a glass electrode that had been rigorously calibrated. The open-circuit potential was subsequently measured with high accuracy using pH meter. Calibration was considered successful, and the reference electrode was deemed ready for use only if the deviation between the measured potential and the theoretical reference value (relative to the standard hydrogen electrode at 25 °C with saturated KCl, SCE: +244 mV, Ag/AgCl: +199 mV) was less than 5 mV.

Electrochemical tests in an H-cell
The electrochemical tests in the H-cell were conducted using a three-electrode system, which was separated by a treated Nafion 117 membrane (size: 2×2 cm2, thickness: 183 μm). Before use, Nafion 117 membrane must undergo a pretreatment process: immerse it in a 5% hydrogen peroxide solution at 80 °C for 1 h; rinse it thoroughly with deionized water for 30 min; then soak it in a 5% sulfuric acid solution at 80 °C for 1 h; finally, wash it again with deionized water for 30 min. The catalyst (2 mg) was dispersed in 360 μL of mixture water and ethanol solution (Vwater:Vethanol = 1:1), along with 40 μL of Nafion (5 wt% aqueous solution), through sonication for 20 minutes to form a homogeneous ink. Then 100 μL of the catalyst ink was loaded onto carbon paper and naturally dried to obtain the working electrode. The geometric area of the working electrode was 1×1 cm2, with a catalyst loading of 0.5 mg·cm-2. The anode electrolyte was 0.5 M KHCO3 (35 mL), and a platinum plate electrode served as the counter electrode, while the cathode electrolyte (35 mL) for coupling reactions was CO2-saturated 0.5 M KHCO3 with 2000 ppm KNO3 (Ar-saturated electrolyte was composed of 0.5 M K2SO4 with 2000 ppm KNO3). Hg/Hg2Cl2 electrode containing saturated KCl solution was used as the reference electrode. To convert applied potentials to RHE reference scale, ERHE = ESCE + 0.0591×pH + 0.241 formula was utilized, where the pH value was set to 7.4. Before conducting electrochemical tests, the cathode electrolyte was purged with CO2 for 30 min. Subsequently, linear sweep voltammetry (LSV) measurements were performed in CO2-saturated electrolyte with KNO3 at a sweep rate of 10 mV/s within the potential range of -0.5 V to -1.5 V (vs. SCE) for 20 min to ensure complete activation of the catalysts. Throughout all electrochemical tests, constant magnetic stirring and CO2 flow rate were maintained.

Urea electrocatalytic synthesis was carried out via controlled potential electrolysis at various potentials for 1 h. The stability test involved continuous controlled potential electrolysis at an applied potential over multiple cycles, with each cycle lasting for 1 hour and refreshing the electrolyte at the beginning of each cycle. Additionally, isotope labeling experiments utilized K15NO3 as a replacement for K14NO3. The electrochemical active surface areas (ECSA) were estimated by calculating the double-layer capacitances (Cdl) from cyclic voltammograms (CV) at scan rates ranging from 40−140 mV/s in a solution of 0.5 M KHCO3 with 2000 ppm KNO3. The differences in current density (Δj = ja - jc) were plotted as a function of the scan rates, and it was determined that the linear slope corresponds to twice the value of Cdl. ECSAs were calculated using the following equation: ECSA = Cdl/Cs, where Cs was the smooth plane capacitance for metal surface (40 μF cm-2)36. The turnover frequency (TOF) of catalysts were calculated using the following formula: TOF = jurea/(16Fn), where jurea was the current density for urea synthesis (A cm-2), F was the Faraday constant (96485 C mol-1), and n was the number of moles of active sites (mol cm-2). CV measurements were conducted at a scan rate of 50 mV/s within the potential range of 0-0.5 V (vs. RHE) in a 0.5 M CO2-saturated KHCO3 solution containing 2000 ppm KNO3. The value of n was determined using: n = Qs/F, where Qs was the surface charge density (C cm-2). By integrating the charge of CV curves, Qs was obtained as half of the integrated value.

Electrochemical tests in a three-electrode flow cell
For the electrochemical tests in a three-electrode flow cell separated by an anion exchange membrane (size: 2×2 cm2, thickness: 100 μm), Ag67Cu33, Nickel foam, Ag/AgCl filled with saturated KCl were employed as the working, counter, and reference electrodes, respectively. Notably, the anion exchange membrane must be subjected to a pretreatment process, which involves sequentially immersing it in 1 M HCl for 1 h, followed by rinsing with deionized water until neutrality. Subsequently, it should be immersed in 1 M NaOH for 1 h, followed by a thorough rinse with deionized water. 4 mg of catalyst was dispersed in a mixture of 360 μL water and ethanol solution (Vwater:Vethanol = 1:4), along with 40 μL of Nafion (5 wt% aqueous solution). The working electrode (1×1 cm2), loaded with 1 mg·cm-2 was prepared following the same procedure as before. The anode electrolyte remained as 0.5 M KHCO3, but the cathode consisted of 0.5 M CO2-saturated KHCO3 with 0.5 M KNO3 (Ar-saturated electrolyte comprised of 0.5 M K2SO4 with 0.5 M KNO3). And electrolytes (35 mL) were continuously injected into their respective compartments during the electrocatalytic process using a peristaltic pump. The applied potentials were converted to the RHE reference scale by ERHE = EAg/AgCl + 0.0591×pH + 0.197, where the pH value was 7.4. LSV measurements and controlled potential electrolysis remained consistent with those used in the H-cell. The long-term stability test consisted of continuous controlled-potential electrolysis under an applied potential for specific cycles; each cycle lasted 3 h and involved refreshing the electrolyte at the beginning of the cycle.

Electrochemical tests in a two-electrode flow cell
For the FOR/urea synthesis systems, Ag67Cu33 was utilized as the working electrode in both the cathode and anode of the two-electrode flow cell. Subsequently, 35 mL of 1 M KOH (with 0.6 M HCOH) and 35 ml of 0.5 M CO2-saturated KHCO3 (with 0.5 M KNO3) were introduced into the anode and cathode compartments at a flow rate of 16 rpm min-1 by a peristaltic pump, respectively, with separation provided by an anion exchange membrane (size: 2×2 cm2, thickness: 100 μm). A catalyst dispersion consisting of 4 mg dispersed in a mixture solution of water and ethanol (Vwater:Vethanol = 1:4), along with 40 μL of Nafion (5 wt% aqueous solution) was prepared. The working electrode (1×1 cm2) loaded with 0.5 mg·cm-2 was obtained using the same procedure as before. Subsequently, both the cathode and anode electrodes required full activation through LSV measurement in the H-cell using previously mentioned methods. Then LSV curves under different conditions in the two-electrode flow cell was measured from 0 V to 2 V at a sweep rate of 50 mV/s. The performance of Ag67Cu33 as a dual-function catalyst was evaluated through constant-current electrolysis at various current densities for 30 min. The cycling tests were conducted at 40 mA cm-2 by ten continuous electrocatalytic experiments, where each cycle lasted 1 h. Subsequently, a long-term stability test was carried out for 60 h at 40 mA cm-2, with the electrolyte being refreshed every 10 h. In all electrochemical tests, after each electrolysis cycle, the electrolytes in the anode and cathode compartments were completely removed via pumping and subsequently replaced with fresh electrolytes before initiating the next cycle.

In-situ Raman measurements
In-situ Raman spectroscopy measurements were conducted with a Renishaw inVia Confocal Raman Microscope equipped with a 532 nm laser (2.5 mW) and an 1800 lines/mm grating, using an Olympus 50X long working distance objective lens (N.A. = 0.5). The electrochemical cell (Supplementary Fig. 32) employed Ag67Cu33 (loaded on 1×1 cm2 carbon paper) as the working electrode, with Pt wire and Ag/AgCl serving as counter and reference electrodes, respectively. A CHI660E electrochemical workstation controlled the system under CO2-saturated 0.5 M KHCO3 with 2000 ppm KNO3, with continuous injection of CO2 during the process. Initially, to simulate the activation process, time-dependent Raman spectroscopy measurements were conducted for 20 min using Linear Sweep Voltammetry (LSV). Subsequently, programmed applied potentials were applied for 300 seconds. Simultaneous spectral acquisition (100−1800 cm-1 range) was performed with 15 accumulations per spectrum to monitor any changes occurring within the catalyst material.

In-situ SR-FTIR measurements
In-situ SR-FTIR measurements were performed at National Synchrotron on Radiation Laboratory (NSRL) BL01B beamline using a test reflector with ZnSe window. The equipment setup integrated a KBr beamsplitter-equipped FTIR spectrometer with liquid nitrogen-cooled MCT detectors, coupled to a Bruker Hyperion 2000 microscope with a ×15 magnifying objective lens. Initially, a micrometer-thick gap between the catalyst electrode and the ZnSe window allowed for reflection observation mode to precisely locate the sample’s surface. Subsequently, measurements were performed using vertically incident infrared light in reflection mode. Spectra were acquired with 128-scan averaging at 4 cm-1 resolution, when the source’s beam current exceeded 400 mA.

Background spectrum was first recorded at open circuit potential (OCP). Electrocatalytic measurements applied sequential potentials from -0.22 to -0.62 V vs. RHE (0.1 V increments) and maintained for 200 s. Importantly, the instrument would automatically subtract OCP-derived backgrounds from test results.

Product quantification
The liquid product analysis employed UV-vis spectrophotometry and 1H NMR spectroscopy, while gas-phase quantification was conducted via GC. The specific detection methods are detailed as follows:

Determination of urea
The urea products were determined using the diacetyl monoxime method. Firstly, 60 mL of concentrated sulfuric acid and 20 mL of concentrated phosphoric acid were slowly added to 120 mL of deionized water. Then 33 mg of FeCl3·6H2O was added and the solution was stirred to form a homogeneous solution (Solution A). Secondly, 500 mg of diacetyl monooxime (DAMO) and 10 mg of thiosemicarbazide (MTSC) were dissolved in 100 mL of deionized water to create a uniform solution (Solution B). Thirdly, 950 μL of deionized water, 50 μL of electrolyte, 2 mL of Solution A and 1 mL of Solution B were sequentially added into a glass bottle. The covered glass bottle containing the mixed solution was then heated to 100 °C for 30 min. After cooling down to 25 °C, the absorbance at 525 nm was measured. Notably, the electrolytes in the H-cell were measured directly, whereas the electrolytes in the flow cells were analyzed following appropriate dilution (50 μL of electrolyte in 950 μL of deionized water). The principle and standard curves of diacetyl monoxime method were shown in Supplementary Fig. 25.

The equation of urea yield rate and Faradic efficiency:
Yield rate (mmol h-1 g-1 cat.) = (Curea (μg mL-1) × V(mL)) / (60.06 × t(h) × m(mg))
FE (%) = (16 × F × Curea (μg mL-1) × V(mL)) / (60.06 × Q) × 100
where Curea = generated urea concentration, V=electrolyte volume, t=reaction time, m=catalyst mass, F=Faraday constant (96485.34) and Q = applied electric quantity.

Determination of ammonia (NH3)
NH3 concentration was determined using Nessler’s reagent spectrophotometry. Briefly, 1 g of potassium sodium tartrate (C4H4KNaO6) was dissolved in 2 mL of deionized water to form homogeneous solution. Then 4.9 mL of deionized water, 100 μL of electrolyte, 0.1 mL of potassium sodium tartrate solution and 0.1 mL of Nessler reagent were added to a glass bottle, respectively. Keeping the mixture at 25 °C for 20 min, its absorbance was then measured at a wavelength of 420 nm, as depicted in Supplementary Fig. 27. The equation for the Faradic efficiency of NH3 was as follows:
FE (%) = (8 × F × CNH3 (μg mL-1) × V(mL)) / (17 × Q) × 100
where CNH3 = generated NH3 concentration, V=electrolyte volume, F=Faraday constant (96485.34) and Q=applied electric quantity.

Determination of NO2--N
Nitrite (NO2-) was detected via N-(1-naphthyl)-ethylenediamine dihydrochloride (NETH) spectrophotometric method. The chromogenic reagent was prepared by dissolving 0.4 g of sulfanilamide, 0.02 g of NETH and 1 mL phosphoric acid (ρ=1.70 g·mL-1) in 5 mL of deionized water. Then a mixture comprising 4.9 mL of deionized water, 100 μL of electrolyte, 0.1 mL of chromogenic reagent was added in a glass bottle, and kept at 25 °C for 20 min. The UV-Vis absorption spectra were obtained at a wavelength of 540 nm, as shown in Supplementary Fig. 28. The equation for the Faradic efficiency of NO2-:
FE NO2- (%) = (2 × F × CNO2- (μg mL-1) × V(mL)) / (46 × Q) × 100
where CNO2- = generated NO2- concentration, V=electrolyte volume, F=Faraday constant (96485.34) and Q=applied electric quantity.

Determination of HCOOH and CH3OH
The identification and quantification of formic acid (HCOOH) and methyl alcohol (CH3OH) were performed using 1H NMR spectroscopy. A mixture containing 300 μL of electrolyte, 200 μL of D2O and 100 μL of DMSO (used as an internal standard) was thoroughly blended and subsequently analyzed using the Ascend 400 NMR spectrometer with accumulating 16 scans, as shown in Supplementary Fig. 30. In view of the disproportionation reaction of HCOH (Cannizzaro Reaction, 2HCHO + OH- → CH3OH + HCOO-) occurring in the presence of OH-, which resulted in the spontaneous formation of HCOOH, the concentration of HCOOH generated through the Ag67Cu33-catalyzed FOR at the anode was calculated using the following formula:
CHCOOH (mmol L-1) = CHCOOH-t - (CCH3OH-t - CCH3OH-0)
where CHCOOH-t and CCH3OH-t represented the concentrations of HCOOH and CH3OH in the post-electrolyte as determined by 1H NMR analysis, and CCH3OH-0 denoted the initial concentration of CH3OH in a 0.6 M HCOH solution (without any treatment). By quantifying the amount of CH3OH generated, we could deduce the quantity of HCOOH formed via the spontaneous decomposition of HCOH based on the chemical equation, which allowed for the accurate calculation of the HCOOH produced via the FOR process catalyzed by Ag67Cu33. Notably, the concentration of CH3OH to inhibit the spontaneous polymerization of HCOH, which must be measured beforehand.

Subsequently, the yield rate and Faradaic efficiency of HCOOH achieved using the Ag67Cu33 catalyst via the FOR process were calculated based on the following equations:
Yield rate HCOOH (mmol h-1 mg-1 cat.) = (CHCOOH (mmol L-1) × V(L)) / (t(h) × m(mg))
FE (%) = (1 × F × CHCOOH (mmol L-1) × V(L)) / Q × 100
where CHCOOH represented the concentration of HCOOH produced by Ag67Cu33 catalyst at the anode, V denoted the volume of the electrolyte solution, t indicated the reaction time, m referred to the mass of the catalyst, F Faraday constant (96485.34), and Q signified the charge.

Determination of H2 and CO
Gas-phase products (H2, CO) were analyzed using GC employing high-purity N2 as the carrier gas, with quantification based on calibration-derived conversion factors from standard reference gases (Supplementary Fig. 29). The equation for the Faradic efficiency of H2 and CO:
FE (%) = (2 × F × CH2 (mmol L-1) × V(L)) / Q × 100
FE (%) = (2 × F × CCO (mmol L-1) × V(L)) / Q × 100
where CH2 = generated H2 concentration, CCO = generated CO concentration, V=total volume of the collected gas, F=Faraday constant (96485.34) and Q=applied electric quantity.

1H NMR method and 15N isotope-labeling experiment
The quantification and identification of urea were accomplished via 1H NMR method and 15N isotope-labeling experiment, respectively. For 1H NMR analysis, 400 μL post-test electrolyte (-0.52 V vs. RHE) was mixed with 200 μL dimethylsulfoxide-d6 (internal standard), and analyzed on an Ascend 600 NMR spectrometer (accumulation of 512 scans). Additionally, various concentrations of 14NH2CO14NH2 mixed uniformly with 200 μL dimethylsulfoxide-d6 were used to construct a standard curve for urea quantification. For the 15N isotope-labeling experiment, CO2-saturated electrolyte with 2000 ppm 15KNO3 was utilized, followed by collecting post-test electrolyte at -0.52 V vs. RHE. The isotope-labeled products were detected through the aforementioned 1H NMR method, with 15NH2CO15NH2 solutions serving as isotopic reference.

Computational detail
All calculations in this study were performed using the Vienna Ab Initio Simulation Package (VASP) based on density functional theory (DFT)52. Projector augmented wave (PAW) pseudopotentials and the semi-local generalized gradient approximation (GGA) with the Perdew–Burke–Ernzerhof (PBE) exchange-correlation functional were employed53. An empirical dispersion-corrected DFT method (DFT-D3) was applied to accurately describe weak long-range van der Waals (vdW) interactions54. The kinetic energy cutoff for the plane-wave expansion was set to 500 eV, and the convergence threshold for the self-consistent field (SCF) iteration was set to 10-5 eV. Geometry optimization was performed using the conjugate gradient method, with forces on each atom being less than 0.03 eV Å-1. A 1×1×1 Gamma k-point mesh was used for k-space integration. For the adsorption model, a vacuum region of approximately 25 Å was added between slabs to prevent interaction between periodic images.

For the Cu(111) and Ag(111) models, 4×4 supercells with 4 atomic layers were used and two bottom layers were fixed. To simulate the interface between Ag(111) and Cu(111), a heterojunction structure of AgCu(111) was constructed, consisting of a (3×7) Ag(111) supercell with 4 atomic layers and a (3×8) Cu(111) supercell with 4 atomic layers. After fixing the z-axis to optimize the lattice and atomic coordinates, the resulting configuration was considered the final structure of the AgCu(111) model, which was then used to study its C-N coupling properties.

To investigate the effect of electrode potential on the C-N coupling process, the CP-VASP code55,56 was employed, with an external potential of -0.96 V (vs. SHE) and a pH of 7.4. Using the conversion relationship between RHE and SHE, the potential relative to RHE was calculated to be -0.52 V, which aligned with the voltage used in the experiment. Since the TARGETMU value corresponding to the vacuum level was -4.57, the applied voltage and pH conditions led to setting TARGETMU to -3.61, and FERMI CONVERGE was set to 0.01.

Implicit solvation was treated using VASPsol++57, which incorporated water solvation in a self-consistent continuous model. The relative permittivity was set to 78.4 to simulate the water electrolyte. The relationship between the Fermi level (EF) and the corresponding electrode potential, referenced to the standard hydrogen electrode (SHE) scale U (V vs. SHE), was formulated as follows:
U (V vs SHE) = -4.57 V - EF / e

The grand canonical energy (GCE) was calculated using the formula below:
GCE = E0 + (CNNE - n) * (-eU)
where E0 was the standard VASP output energy, CNNE was the charge-neutral electron number, n was the electron number, and U was the external electrode potential. The grand canonical energy (GCE) was expressed in terms of eV. Note that E0 did not account for changes in the electron number and was no longer an accurate representation of the system’s total energy.

The change in reaction free energy could be obtained using the following equation:
ΔG = ΔGCE + ΔZPE - TΔS + n1*PH*kB T ln 10 + n2 * eU
where ΔGCE was the total grand canonical energy difference before and after intermediate adsorption, ΔZPE represented the difference in zero-point energy, ΔS was the difference in entropy, kB was the Boltzmann constant, and T was the temperature, taken as 298.15 K. n1 was the number of proton transfers involved in the corresponding step, and n2 referred to the number of electron transfers. The pH value used in this study was set to 7.4, and U was -0.96 V. The zero-point energy and entropy of free molecules and adsorbates were obtained from vibrational frequency calculations. Since VASP could not accurately describe charged systems, the energy of NO3- needed to be measured according to this literature58.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "In this research, the authors developed AgCu biphasic aerogel catalysts with an analogous rectifying interface, which effectively facilitated the synthesis of urea via the coupling reaction of CO2 and NO3-, achieving an ultrahigh urea yield rate (54.8 mmol h-1 gcat.-1) with a Faradaic efficiency (36.6%) at a low potential (-0.52 V vs. RHE). In-situ characterization and theoretical calculations revealed that the analogous rectifying interface effect could generate stable Cuδ+ active sites near the interfaces, thereby enhancing the adsorption of *NOH and stabilizing the *CO intermediate, which improved the C-N coupling dynamics between *CO and *NOH. Furthermore, the two-electrode flow cell utilized Ag67Cu33 as a dual-function catalyst, demonstrating remarkable efficacy in the production of both urea and HCOOH, which underscored its practical applicability. The manuscript is well-structured and presents compelling findings regarding urea production. I recommend the acceptance of this manuscript after minor revisions. Below are some comments for the authors:\n1.The authors introduced the concept of “analogous rectifying interface effect” in this study, and chose silver (Ag) and copper (Cu) as the interface materials. Can you provide detailed explanation on this “effect”, and clarify the rationale behind the selection of Ag and Cu for this study?\n2.In this study, the authors synthesized a series of AgCu catalysts and evaluated their performance in urea synthesis, and the electrochemical active surface area (ECSA) was measured to investigate their electrocatalytic activity. Can you provide the turnover frequency (TOF) in addition to Cdl for a more comprehensive analysis of the intrinsic catalytic activity?\n3.In Figure 1d, the authors investigated the Ag-Cu interfaces using HAADF-STEM, revealing two distinct lattice spacings. I recommend that the authors include the HAADF-STEM image intensity profile of these two regions in the Supporting Information to enhance the reliability of the findings.\n4.Within the two-electrode flow cell setup, the authors integrated urea synthesis with the formaldehyde oxidation reaction (FOR). While the authors have provided XRD results of catalysts post-experiment to demonstrate stability, I suggest to supplement the XPS results of catalysts to further substantiate the absence of oxidation in the catalysts.\n5.The DFT results presented in Figure 4e indicated that the catalytic center within the catalyst was Cu. However, the experimental results depicted in Figure 2c showed that urea production initially increased and subsequently decreased with rising Cu content. Can the authors offer a detailed explanation for this observed phenomenon?\n6.The authors have identified all by-products generated during urea synthesis (Figure S8), but the standard curves for CO and H2 are absent, and should be included in the Supporting Information.\n7.In this study, the authors discovered a novel pathway for urea synthesis utilizing the AgCu catalyst (Figure 4e), demonstrating that both key C-N coupling reactions occurred via a spontaneous process, which was advantageous for urea synthesis from a thermodynamic perspective. However, to establish that these reactions are also kinetically favorable, can the authors calculate the reaction pathways for the first and second C-N couplings, such as those reported in Adv. Mater. 2024, 36, 2402160.", "reviewer_2": "In this work, the authors constructed AgCu biphasic aerogels to promote urea synthesis from the co-reduction of NO3- and CO2. The results showed that the proposed exhibit high catalytic performance. Since I am a researcher engaged in theoretical computational work, I will present my question solely from the perspective of density functional theory.\n1) According to a recent similar work (Nature Communications 2024, 15, 10331), the similar AgCu biphasic heterostructure exhibits high catalytic performance for producing multicarbon products. However, the urea synthesis and CO2 reduction may compete with each other, which may affect the selectivity of this catalyst. Therefore, the authors should evaluate the selectivity of this catalyst through theoretical calculations, such as its catalytic activity for hydrogen evolution reactions, nitrate reduction, and CO2 reduction to multi-carbon products.\n2) C-N coupling is a central challenge in urea synthesis, as it is crucial to the kinetics of the reaction. However, the study does not delve deeply enough into this issue. As mentioned in the paper, different nitrogen-containing intermediates, such as NO and NOH*, are generated during the nitrate reduction process, but the author only considers the coupling between NOH* and CO. While this coupling may indeed be experimentally observed, to highlight the superiority of this coupling, the interactions between other nitrogen-containing intermediates and CO or CO2 should also be considered. More importantly, the author mentions that this coupling can occur spontaneously, which may be inaccurate, as this explanation is solely based on thermodynamics. The author has not calculated the kinetic barriers for these reactions, and therefore, this aspect should also be investigated.\n3) The author claims that the formation of AgCu heterojunctions facilitates the reduction of CO2 to CO. However, theoretical calculations show that the Gibbs free energy for this process is as high as 1.27 eV, which is higher than the 0.99 eV for traditional Cu catalysts. Based on this data, the AgCu catalyst does not exhibit good performance for CO formation, as it hinders the formation of CO*, which is detrimental to urea synthesis. This contradicts the experimental results. I suspect this may be related to the CHE model used in the calculations, as this model does not account for the effects of pH and potential, and the author also neglects the influence of solvent effects.\n4) The author considers the impact of pH on catalytic activity in the computational method, with a selected pH value of 7. However, the CHE model oversimplifies the influence of pH on catalytic activity, often yielding results that contradict experimental observations. For detailed discussion, please refer to a series of articles by Zhiyao Duan et al. on calculating under constant potential. Therefore, to better account for the effect of pH, we suggest re-evaluating the results using the constant potential method.\n5) To make readers better understand the AgCu structure, the corresponding contar files computed by VASP should be provided, including catalyst, adsorbed intermediates.", "reviewer_3": "In this article, Feng et al. constructed analogous rectifying interfaces in AgCu biphasic aerogels with stable Cuδ+ active sites, which simultaneously promoted *NOH adsorption and *CO coverage, thereby improving C-N coupling dynamics. Specifically, using Ag67Cu33 in a flow cell, ultrahigh urea yield (54.8 mmol h-1 gcat.-1) with notable Faradaic efficiency (FE, 36.6%) at a low potential (-0.52 V vs. RHE) was achieved. Overall the manuscript is quite impressive. Besides, the manuscript can be revised to address some issues before being considered for publication in this journal.\n1. How do the specific concentrations of NaBH4 and the reaction conditions influence the formation of the AgCu biphasic aerogels? Could variations in these parameters affect reproducibility or performance consistency?\n2. The LSV figure shows that the reduction starts at nearly -0.1V. Why were specific potential ranges (-0.22 V to -0.62 V vs. RHE) selected for electrocatalysis? Could alternative ranges provide insights into different catalytic behaviours?\n3. Have the authors considered conducting control experiments using pure Ag and pure Cu catalysts in physically mixed configurations instead of biphasic aerogels to further validate the role of the affinity-modifying interface? Also, the authors should provide control experiments for urea synthesis using pure Ag and pure Cu catalysts.\n4. Additional surface or bulk characterization techniques can be considered (e.g., XANES, EXAFS) to further confirm the oxidation state of Cu and Ag during the electrocatalytic process.\n5. Stability tests show no significant morphological changes, but what about the possible leaching of Ag or Cu? ICP-MS can be used to confirm the absence of dissolved metals for long-term use.\n6. The authors discuss electrochemical urea synthesis, a promising alternative to the traditional urea synthesis process using bimetallic systems. They should add some discussion from recent literature to keep the manuscript up-to-date using catalysts based on metal-N4 or metal phthalocyanine catalysts.\n7. The pH value of the electrolyte may be varied accompanied with ammonia production (by-product from NO3– reduction). How the catalytic stability can be maintained?\n8. The authors carry out the DAMO method and Nessler's reagent test to detect urea and ammonia if produced from NO3– reduction. However, all the UV data for the samples are missing in the manuscript as well as in the supporting information. The author should add that missing information.\n9. Density functional theory (DFT) models are based on specific crystallographic orientations. How would variations in the crystallographic alignment of the Ag and Cu interfaces affect the results?\n10. How was the 0.5 M KNO3 concentration determined to be optimal for flow cell experiments? Experiments can be conducted at different concentrations."}, "author_response": "We extend our sincere appreciation to the reviewer for their comprehensive evaluation of our work. We are particularly grateful for the invaluable and professional insights they provided, especially regarding the active sites for urea synthesis from both experimental and theoretical perspectives. The reviewer also constructively recommended incorporating supplementary evidence to enhance the reliability of our work. We are committed to providing detailed responses to address the reviewer's questions and firmly believe that the quality and clarity of our manuscript will be significantly enhanced after incorporating these revisions in accordance with your feedback.\nIn response to the reviewer's insightful comments, we have made the following modifications:\n1. Detailed exposition on the “analogous rectifying interface effect” and material selection: We thoroughly outlined our innovative material design philosophy, explored the significant implications of this “effect”, and presented compelling justifications for selecting Ag and Cu.\n2. The turnover frequency (TOF) as an indicator of intrinsic catalytic activity: To provide a more comprehensive assessment of the intrinsic catalytic performance in urea synthesis, we conducted an in-depth analysis of the TOF for each catalyst.\n3. The absence of HAADF-STEM image intensity profiles: In accordance with the reviewer's suggestion, we have supplemented HAADF-STEM and HR-TEM image intensity profiles at the interfacial regions to further substantiate the reliability of the interface formation.\n4. XPS analysis of Ag67Cu33 in the two-electrode system: We have conducted a re-evaluation of the long-term stability in the two-electrode system, employing XPS on catalysts at both cathode and anode after testing to identify any potential changes in their chemical states.\n5. The decreased urea synthesis efficiency with increasing Cu content: We conducted an in-depth analysis of the active sites involved in urea synthesis, and in conjunction with experimental data, elucidated the rationale for the volcano-like trend observed over AgxCu100-x catalysts in urea synthesis.\n6. The lack of standard curves for CO and H2: In response to the reviewer's insightful suggestion, we have included the standard calibration curves for H2 and CO in the Supporting Information.\n7. The calculations of the reaction pathway for the C-N couplings: We conducted DFT calculations to evaluate the kinetic barriers for two critical coupling steps, which suggested that urea production on Ag67Cu33 offered both thermodynamic and kinetic advantages in the C-N coupling processes.\nWe extend our sincerest gratitude for your invaluable feedback and the time you had dedicated to reviewing our work. We hope that our meticulously revised manuscript, along with detailed explanations, will adequately address your concerns and fully meet your expectations. We eagerly anticipate your further insights.\nA comprehensive response to the reviewer's insightful suggestions is provided below.\n1.The authors introduced the concept of “analogous rectifying interface effect” in this study, and chose silver (Ag) and copper (Cu) as the interface materials. Can you provide detailed explanation on this “effect”, and clarify the rationale behind the selection of Ag and Cu for this study?\nAuthor response: We would like to extend our sincere gratitude to the reviewer for raising this pivotal question. As widely acknowledged, copper (Cu) has long been recognized as one of the most promising catalysts for urea synthesis due to its exceptional performance in both carbon dioxide reduction reaction (CO2RR) and nitrate reduction reaction (NO3RR). Extensive prior research had demonstrated that Cuδ+ sites were particularly advantageous active sites for these reactions, as they could modulate the adsorption behavior of intermediates (J. Am. Chem. Soc. 2023, 145, 26133-26143; Adv. Mater. 2024, 36, 2409697). However, maintaining Cuδ+ under applied negative potentials during reduction reactions posed significant challenges, leading to suboptimal catalyst stability. Drawing inspiration from the rectifying interface effect observed in heterojunctions, we hypothesized that constructing a built-in electric field could potentially address this issue. Nevertheless, the regulation of electron redistribution at heterointerfaces was a complex process, influenced by multiple factors including work functions, band structures, and dielectric constants, which were critical in determining the formation of the built-in electric field but were also susceptible to external perturbations. Consequently, we proposed substituting the semiconductor with another metal to achieve an analogous rectifying interface effect. This method facilitated electron transfer across interfaces via the electronegativity differences between metals, offering greater controllability.\nIn consideration of the exceptional properties and facile synthesis method of bimetallic aerogels, we selected Ag and Cu as the constituent materials. Cu was identified as the active sites for urea synthesis due to its established reputation as an ideal catalyst. The selection of Ag as the secondary metal was motivated by several key factors: firstly, its unique capability to form aerogels, a property shared by only approximately 10% of metals in the periodic table, predominantly precious ones (Matter 2021, 4, 54-94); secondly, its thermodynamic immiscibility with Cu facilitated the formation of critical interfaces, which was crucial for the desired effect; thirdly, its higher electronegativity promoted the induction of Cuδ+, thereby enhancing the catalytic activity; fourthly, its superior electrical conductivity and remarkable stability were essential during the electrocatalytic process; finally, its relatively lower cost compared to other potential metals rendered it a pragmatic choice. In summary, by incorporating Ag as the secondary metal, we had successfully constructed analogous rectifying interfaces with Cu, enabling precise control over electron redistribution and stabilizing Cuδ+, thus significantly enhancing the urea synthesis efficiency.\n2.In this study, the authors synthesized a series of AgCu catalysts and evaluated their performance in urea synthesis, and the electrochemical active surface area (ECSA) was measured to investigate their electrocatalytic activity. Can you provide the turnover frequency (TOF) in addition to Cdl for a more comprehensive analysis of the intrinsic catalytic activity?\nAuthor response: We would like to extend our heartfelt thanks to the reviewer for their valuable and professional suggestions. To gain a deeper understanding of the intrinsic catalytic activity of AgxCu100-x catalysts, we evaluated the turnover frequency (TOF) of each catalyst. The TOF signified the rate of electron consumption per unit time at an individual active site (Angew. Chem. Int. Ed. 2024, 64, e202413534) and was determined using the following formula: TOF = jurea / (16Fn), where jurea was the current density for urea synthesis (A cm-2), F was the Faraday constant (96485 C mol-1), and n was the number of moles of active sites (mol cm-2). Additionally, n was determined via the formula: n = Qs / F, where Qs was the surface charge density (C cm-2). To obtain these values, cyclic voltammetry (CV) measurements were conducted at a scan rate of 50 mV/s within the potential range of 0-0.5 V (vs. RHE) in a 0.5 M CO2-saturated KHCO3 solution containing 2000 ppm KNO3. By integrating the charge of CV curves, we obtained the Qs as half of the integrated value, from which we derived the values of n and subsequently the TOF for each catalyst.\nAs illustrated in Figure S15, the calculated Qs values for Ag, Ag80Cu20, Ag67Cu33, Ag50Cu50, and Cu were 3.95 × 10-5, 1.033 × 10-4, 1.17 × 10-4, 9.97 × 10-5, and 6.38 × 10-5 C cm-2, respectively. Subsequently, the corresponding n values were determined to be 4.09 × 10-10, 1.07 × 10-9, 1.21 × 10-9, 1.03 × 10-9, and 6.61 × 10-10, indicating that Ag67Cu33 possessed the highest number of active sites under identical catalyst loading conditions, which was consistent with the ECSA results. Furthermore, the TOF values of various AgxCu100-x catalysts were compared in Figure S15f, revealing that Ag67Cu33 exhibited the largest TOF value across the entire applied potential range. In conjunction with the ECSA findings, both results confirmed that Ag67Cu33 demonstrated superior intrinsic catalytic activity, thus enhancing urea synthesis performance. Please refer to line 218 on page 7 of the revised manuscript, where further details are provided.\nFigure S15. CV curves of Ag(a), Ag80Cu20 (b), Ag67Cu33 (c), Ag50Cu50 (d) and Cu(e) within the potential range of 0-0.5 V (vs. RHE). (f) Comparative analysis of the TOF for urea synthesis over various AgxCu100-x catalysts.\n3.In Figure 1d, the authors investigated the Ag-Cu interfaces using HAADF-STEM, revealing two distinct lattice spacings. I recommend that the authors include the HAADF-STEM image intensity profile of these two regions in the Supporting Information to enhance the reliability of the findings.\nAuthor response: We are deeply appreciative of the reviewer's insightful comments regarding the omission of crucial evidence concerning the Ag-Cu interfaces. In response to the reviewer's valuable feedback, we had incorporated the HAADF-STEM image intensity profiles for these two regions in Supplementary Figure S4, which revealed the two different lattice spacings were 0.237 nm and 0.208 nm, corresponding precisely to Ag(111) and Cu(111), respectively. Furthermore, we had additionally supplemented the HR-TEM image intensity profiles around the Ag-Cu interface regions of Ag67Cu33 both before and after electrochemical measurements (Figure S3 and S11), which were consistent with HAADF-STEM results, and significantly enhance the robustness and reliability of our findings.\nFigure S4. HAADF-STEM image(c) of Ag67Cu33, along with the corresponding intensity profiles of the Ag regions(a-b) and Cu regions(d-e) in designated areas.\nFigure 3. HR-TEM image(a), EDX mapping(b) and the corresponding concentration result(c) of Ag67Cu33. Intensity profiles of the Ag regions(d) and Cu regions(e) within the white-boxed areas adjacent to the Ag-Cu interface(a).\nFigure 11. Characterizations of Ag67Cu33 after stability test. TEM image(a), HR-TEM image(b) and corresponding intensity profiles of the Ag regions(c) and Cu regions(f) within the white-boxed areas(b), high resolution XPS spectra of Ag 3d(d) and Cu 2p(e).\n4.Within the two-electrode flow cell setup, the authors integrated urea synthesis with the formaldehyde oxidation reaction (FOR). While the authors have provided XRD results of catalysts post-experiment to demonstrate stability, I suggest to supplement the XPS results of catalysts to further substantiate the absence of oxidation in the catalysts.\nAuthor response: We express our sincere gratitude for your invaluable feedback regarding the missing critical XPS analysis after long-term stability tests, which was essential for elucidating the chemical states of the Ag67Cu33 catalyst. In response to the reviewer's recommendations, we had re-evaluated the stability in a two-electrode system and analyzed the results using XPS. As shown in Figure S58, the high-resolution XPS spectra of Ag 3d for Ag67Cu33 revealed peaks at 368.28 eV and 374.31 eV at the cathode (d), and 368.27 eV and 374.29 eV at the anode (f), corresponding to the Ag 3d5/2 and Ag 3d3/2 of Ag0. Similarly, the Cu 2p spectra exhibited peaks at 932.60 eV and 952.48 eV at the cathode (e), and 932.65 eV and 952.52 eV at the anode (g), attributed to Cu 2p3/2 and Cu 2p1/2 of Cu0/Cu+. Although minor peaks indicative of Cu2+ were present, the surface oxidation observed in the ex-situ XPS measurements could be attributed to exposure to air. These findings were consistent with the pre-test XPS data of Ag67Cu33 (Figure 1g and 1h), corroborated by SEM images and XRD patterns, thereby confirming the integrity and stability of the catalyst's morphology and structure. Please refer to line 446 on page 13 of the revised manuscript.\nFigure S58. The SEM images of Ag67Cu33 at the cathode(a) and anode(b), along with XRD results(c), and corresponding high resolution XPS spectra of Ag 3d and Cu 2p at the cathode(d-e) and anode(f-g) in the two-electrode system after long-term electrolysis for 60 h.\n5.The DFT results presented in Figure 4e indicated that the catalytic center within the catalyst was Cu. However, the experimental results depicted in Figure 2c showed that urea production initially increased and subsequently decreased with rising Cu content. Can the authors offer a detailed explanation for this observed phenomenon?\nAuthor response: We extend our sincere gratitude to the reviewer for highlighting this intriguing question. We had synthesized a series of AgxCu100-x catalysts and evaluated their performance in urea synthesis, noting significant variations (Figure 2c). According to our density functional theory (DFT) calculations, Cuδ+ sites adjacent to the interfaces served as the active centers for urea synthesis, supported by previous studies. For CO2RR, research on Cu+-rich ultrathin two-dimensional Cu2O nanosheets had demonstrated that Cu+ enhanced the adsorption capacity of *CO (J. Am. Chem. Soc. 2023, 145, 26133-26143). In NO3RR, it had also been observed that Cuδ+ facilitated rapid reduction from NO3- to NH3 (Appl. Catal. B Environ.317, 121811 (2022)). More importantly, recent studies on urea synthesis had confirmed that stable electron-deficient Cu played a pivotal role in enhancing urea synthesis efficiency (Angew. Chem. Int. Ed. 2024, 63, e202410938; Adv. Mater. 2024, 36, 2409697).\nTherefore, we proposed that the initial increase and subsequent decrease in urea synthesis performance with rising Cu content could be attributed to the fact that the number of Ag-Cu interfaces did not proportionally increase with higher Cu content. This resulted in the actual active sites (Cuδ+) for urea synthesis not genuinely increasing but rather the number of Cu0 sites increasing, leading to more by-products as shown in Figure S9. Consequently, the volcano-like trend in urea synthesis performance for AgxCu100-x could be ascribed to fluctuations in the number of Cuδ+ active sites.\nFurthermore, the ECSA (Figure S14) and TOF (Figure S15) results demonstrated that Ag67Cu33, which exhibited the highest performance, also possessed superior intrinsic catalytic activity for urea synthesis, thereby further validating our conclusion.\n6.The authors have identified all by-products generated during urea synthesis (Figure S8), but the standard curves for CO and H2 are absent, and should be included in the Supporting Information.\nAuthor response: We extend our sincere gratitude for your timely reminder. Following your insightful suggestion, we have now incorporated the standard calibration curves of H2 and CO into Figure S28, which significantly contributed to the accurate quantification of the products.\nFigure S28. Standard curves for quantification of H2 (a), CO (b).\n7.In this study, the authors discovered a novel pathway for urea synthesis utilizing the AgCu catalyst (Figure 4e), demonstrating that both key C-N coupling reactions occurred via a spontaneous process, which was advantageous for urea synthesis from a thermodynamic perspective. However, to establish that these reactions are also kinetically favorable, can the authors calculate the reaction pathways for the first and second C-N couplings, such as those reported in Adv. Mater. 2024, 36, 2402160.\nAuthor response: We sincerely appreciate the reviewer's insightful suggestions regarding the in-depth investigation of the two crucial C-N coupling processes. In response, we performed DFT calculations (as detailed in Computational detail, page 18 of the revised manuscript) to determine the kinetic barriers for these coupling steps (Adv. Mater. 2024, 36, 2402160; Angew. Chem. Int. Ed. 2023, 62, e202305447). As shown in Figure S35a, the first C-N coupling step involving *NOH and *CO exhibited a transition state (TS) energy barrier of 0.67 eV, accompanied by a thermodynamically spontaneous process with a free energy change of ΔG = -0.63 eV. Subsequently, the second C-N coupling from *CONH2 and *NO demonstrated a TS energy barrier of 0.87 eV (Figure S35c) with an energetically favorable step (ΔG = -0.38 eV). These results indicated that the critical C-N coupling processes on the AgCu catalyst required relatively low activation barriers with exothermic steps, suggesting that the urea formation pathway benefited from both thermodynamic feasibility and kinetic accessibility in the C-N coupling steps. Please refer to line 351 on page 11 of the revised manuscript.\nFigure S35. The reaction pathways and their corresponding potential-time of the system during the slow-growth process under constant potential method: (a-b) the first C-N coupling (forming *CONOH), (c-d) the protonation of *NOH (yielding *CO + *NHOH), (e-f) the second C-N coupling (forming *NOCONH2), and (g-h) the protonation of *NO (yielding *NOH + *CONH2). The insert was the structures corresponding to the initial, transition, and final states are illustrated. The blue, red, yellow, green, gray and white balls represent Cu, Ag, O, N, C, H atoms, respectively.\nReviewer #2:\nIn this work, the authors constructed AgCu biphasic aerogels to promote urea synthesis from the co-reduction of NO3- and CO2. The results showed that the proposed exhibit high catalytic performance. Since I am a researcher engaged in theoretical computational work, I will present my question solely from the perspective of density functional theory.\nAuthor response: We express our sincere gratitude to the reviewer for the rigorous evaluation of our work and provision of expert insights. These constructive recommendations have not only enhanced manuscript quality through error correction, but also significantly advanced our understanding of computational methodologies in electrocatalytic urea synthesis. In strict accordance with the reviewer's professional guidance, we have implemented revised computational methods and expanded details in the revised manuscript, hoping our revisions are in accordance with the reviewer's expectations. The comprehensive DFT calculation data and necessary revised calculation method are presented below:\n1) We initially conducted experiments to verify the absence of C2+ products using our Ag67Cu33 catalyst, followed by DFT calculations investigating HER, NO3RR, and CO2-to-C2+ pathways to elucidate the factors affecting the selectivity of the catalyst.\n2) We conducted a comprehensive analysis of the coupling trends among all potential reactions in the urea synthesis pathway, while analyzing kinetic barriers for two critical C-N couplings to elucidate their formation propensity from the kinetic perspective.\n3) We accounted for pH, potential, and solvent effects to refine our DFT calculation methodology, followed by a thorough re-evaluation of the CO2-to-CO pathway on Ag, Cu and Ag67Cu33, in order to assess their *CO formation capabilities.\n4) We systematically recalibrated all theoretical findings using the constant potential method, establishing pH and potential to align with actual experimental conditions, thereby improving the reliability of the derived results.\n5) We provided top and side views of DFT-optimized structures for enhanced readers' understanding, and all VASP-generated CONTCAR files have been submitted in accordance with the data policies of Nature Communications.\nBased on reviewer's expert feedback, we have revised the calculation method as detailed below:\nAll calculations in this study were performed using the Vienna Ab Initio Simulation Package (VASP) based on density functional theory (DFT) (J. Comput. Chem. 2008, 29, 2044-2078). Projector augmented wave (PAW) pseudopotentials and the semi-local generalized gradient approximation (GGA) with the Perdew–Burke–Ernzerhof (PBE) exchange-correlation functional were employed (Phys. Rev. Lett. 1996, 77, 3865-3868). An empirical dispersion-corrected DFT method (DFT-D3) was applied to accurately describe weak long-range van der Waals (vdW) interactions (J. Chem. Phys. 2010, 132). The kinetic energy cutoff for the plane-wave expansion was set to 500 eV, and the convergence threshold for the self-consistent field (SCF) iteration was set to 10-5 eV. Geometry optimization was performed using the conjugate gradient method, with forces on each atom being less than 0.03 eV Å-1. A 1 × 1 × 1 Gamma k-point mesh was used for k-space integration. For the adsorption model, a vacuum region of approximately 25 Å was added between slabs to prevent interaction between periodic images.\nFor the Cu(111) and Ag(111) models, 4 × 4 supercells with 4 atomic layers were used and two bottom layers were fixed. To simulate the interface between Ag(111) and Cu(111), a heterojunction structure of AgCu(111) was constructed, consisting of a (3 × 7) Ag(111) supercell with 4 atomic layers and a (3 × 8) Cu(111) supercell with 4 atomic layers. After fixing the z-axis to optimize the lattice and atomic coordinates, the resulting configuration was considered the final structure of the AgCu(111) model, which was then used to study its C-N coupling properties.\nTo investigate the effect of electrode potential on the C-N coupling process, the CP-VASP code (J. Am. Chem. Soc. 2021, 143, 9423-9428; J. Am. Chem. Soc. 2023, 145, 25352-25356) was employed, with an external potential of -0.96 V (vs. SHE) and a pH of 7.4. Using the conversion relationship between RHE and SHE, the potential relative to RHE was calculated to be -0.52 V, which aligned with the voltage used in the experiment. Since the TARGETMU value corresponding to the vacuum level was -4.57, the applied voltage and pH conditions led to setting TARGETMU to -3.61, and FERMICONVERGE was set to 0.01.\nImplicit solvation was treated using VASPsol++ (J. Chem. Phys. 2023, 159), which incorporated water solvation in a self-consistent continuous model. The relative permittivity was set to 78.4 to simulate the water electrolyte. The relationship between the Fermi level (EF) and the corresponding electrode potential, referenced to the standard hydrogen electrode (SHE) scale U (V vs. SHE), was formulated as follows: The grand canonical energy (GCE) was calculated using the formula below:\nwhere E0 was the standard output energy, N was the charge-neutral electron number, n was the electron number, and U was the external electrode potential. The grand canonical energy (GCE) was expressed in terms of eV. Note that E0 did not account for changes in the electron number and was no longer an accurate representation of the system's total energy.\nThe change in reaction free energy could be obtained using the following equation:\nwhere ΔG was the total grand canonical energy difference before and after intermediate adsorption, ΔZPE represented the difference in zero-point energy, ΔS was the difference in entropy, kB was the Boltzmann constant, and T was the temperature, taken as 298.15 K. n1 was the number of proton transfers involved in the corresponding step, and n2 referred to the number of electron transfers. The pH value used in this study was set to 7.4, and USHE was -0.96 V. The zero-point energy and entropy of free molecules and adsorbates were obtained from vibrational frequency calculations. Since VASP could not accurately describe charged systems, the energy of NO3- needed to be measured according to this literature (Adv. Funct. Mater. 2021, 31, 2008533).\nWe would like to express our heartfelt gratitude once again to the reviewer for investing their valuable time and effort in evaluating our manuscript. The professional and insightful feedback we received has significantly enhanced the quality of our work, providing invaluable guidance for our future researches on C-N coupling. In alignment with reviewer's suggestions, we have refined and optimized our calculation method, recalibrated all of computational data in the article, and incorporated additional details to improve the overall quality of our manuscript. We trust our revised manuscript meets your expectations and look forward to your further feedback. The detailed response is outlined below:\n1) According to a recent similar work (Nature Communications 2024, 15, 10331), the similar AgCu biphasic heterostructure exhibits high catalytic performance for producing multicarbon products. However, the urea synthesis and CO2 reduction may compete with each other, which may affect the selectivity of this catalyst. Therefore, the authors should evaluate the selectivity of this catalyst through theoretical calculations, such as its catalytic activity for hydrogen evolution reactions, nitrate reduction, and CO2 reduction to multi-carbon products.\nAuthor response: We would like to extend our heartfelt gratitude to the reviewer for raising this critical question. It is widely acknowledged that parallel competitive reactions, such as the carbon dioxide reduction reaction (CO2RR), nitrate reduction reaction (NO3RR), and hydrogen evolution reaction (HER) would significantly compromise the efficiency of urea synthesis. While Zhang et al. successfully synthesized an Ag-Cu biphasic heterostructure exhibiting remarkable catalytic performance towards C2+ products (Nat. Commun.2024,15, 10331), it was noteworthy that their required applied potential was more negative than -0.7 V vs. RHE. In contrast, the potential range utilized in our work spanned from -0.22 V to -0.62 V vs. RHE. The different potential window may result in the coupling of N- and C-intermediates or their undergoing further hydrogenation independently (Nat. Commun. 2024, 15, 1095). Furthermore, their research also demonstrated that the Ag:Cu ratio exerted a significant influence on the predominant products, with an observed shift from C2+ compounds to CO as the proportion of Ag increases. To investigate the selectivity of the catalyst, we first evaluated the performance of Ag67Cu33 towards CO2RR using a CO2-saturated KHCO3 electrolyte, which indicated that only CO and H2 were produced, confirming the absence of any C2+ products in our work (Figure S16). Furthermore, the current density observed during the co-reduction of CO2 and NO3- was significantly higher than that achieved in their individual reduction processes, suggesting a potential coupling process.\nTo gain a more comprehensive understanding of the complex factors, we conducted DFT calculations on HER, NO3RR, and CO2-to-C2+ processes over the Ag(111)-Cu(111) (Ag67Cu33) aiming to elucidate these in greater detail. Initially, we investigated the synthesis of urea on Ag67Cu33 (Figure S34) from the adsorption of CO2 and NO3-, while also examining their respective hydrogenation pathways in CO2RR and NO3RR, which demonstrated that C-N coupling occurred preferentially, leading to the urea formation, with the rate-determining step (RDS) being the adsorption of NO3-. Notably, the crucial C-N couplings were not only thermodynamically spontaneous but also kinetically favorable. As depicted in Figure S46a, Ag67Cu33 exhibited a preference for *NOH intermediate in NO3RR, where *CO + *NOH formed critical *CONOH intermediate during urea synthesis. Specifically, after CO2 adsorption, both the thermodynamic energy barrier (Figure S34) and the transition state (TS) energy barrier (Figure 35) for C-N coupling (*NOH + *CO → *CONOH) were lower than those for the protonation of *NOH (*NOH + *CO → *NHOH + *CO), which favored the selective occurrence of C-N coupling over NO3RR. For HER (Figure S46b), the RDS involved the desorption of H2 with an energy barrier of 0.94 eV, which was higher than that of the RDS in C-N coupling (0.88 eV), thereby reducing the competitiveness of HER process.\nNext, we explored the directed hydrogenation and C-C coupling within the CO2-to-C2+ process (Figure S49). It was evident that the transformation of CO2 + CO2 into *CO + *CO was favored over other protonation pathways and C-C couplings. Notably, the subsequent step required an energy input of 0.28 eV to form *CO + *CHO, and a significantly higher free energy barrier of 0.89 eV to generate *COCO (highlighted in pale pink). However, *CO also emerged as the critical C-intermediate for the first C-N coupling, which proceeded downhill with a low TS energy barrier (Figure S34), which rendered it more inclined toward C-N coupling rather than the protonation of *CO. Following multistep hydrogenation, key intermediates such as *COCH2 and *CHOCH2 were formed, facilitating the production of C2+ products like C2H4 and CH3CH2OH (Nat. Commun.2024,15, 10331; Nature Energy 2024, 9, 1485-1496), while the formation of *CO + *CH4 still remained more facile (indicated in light blue). Although the production of C2+ and hydrocarbons was theoretically feasible, it necessitated multiple hydrogenation and deep reductions, requiring potentials (Angew. Chem. Int. Ed., e202503417;Appl. Catal. B Environ. 2023, 325, 122310) far more negative than those applied in our work. In contrast, C-N coupling occurred more readily which utilized *CO as the reaction intermediate. Furthermore, considering the actual process, NO3- exhibited exceptional water solubility as free ions, while the low solubility of CO2 in water introduced resistance during its transition from the gas phase to the liquid phase and subsequent diffusion to the electrode surface (Chem. Eng. J. 2021, 405, 126975). Although the negatively charged NO3- ions experienced increased electrostatic repulsion against the negatively charged electrode, which made their adsorption slightly more challenging (Angew. Chem. Int. Ed. 2024, 63, e202402684), their faster diffusion rate compared to CO2 allowed them to effectively occupy sufficient active sites. Consequently, N-intermediates and C-intermediates could intermingle, thus promoting C-N coupling.\nIn conclusion, during urea synthesis, HER, NO3RR, and CO2-to-C2+ processes could occur concurrently. Cuδ+ active sites were capable of generating NH3 or C2+ products upon the introduction of either NO3- or CO2 (Nat. Commun. 2023, 14, 7383). However, when CO2 and NO3- were introduced simultaneously, C-N coupling became preferential due to several critical factors. For NO3RR, the formation of *NOH was facilitated, and the selection of *NOH + *CO to form *CONOH was more favorable than *NHOH + *CO from both thermodynamic and kinetic perspectives. Regarding HER, the free energy barrier of RDS surpassed that of urea synthesis, thereby suppressing H2 production. For CO2-to-C2+, *CO + *CO demanded considerably higher energy than C-N coupling to enable subsequent protonation or C-C coupling processes, and the synthesis of C2+ species or hydrocarbons involved multiple hydrogenation steps and deep reductions, requiring a far more negative electrochemical potential. Macroscopically, two factors might modulate its catalytic behaviors. First, an optimal potential range could not only enhance protonation but also ensure that the coupling process remained dominant over side reactions. However, excessively negative potentials significantly increased the occurrence of undesired side reactions. Second, the factors influenced the mass transfer rates of CO2 and NO3- (such as the flow rate of CO2 or the concentration of NO3-), which in turn affected the effective adsorption of C-/N-intermediates at active sites. Elevated concentrations of either CO2 or NO3- might competitively occupy more active sites, thereby increasing the side reactions. Please refer to line 391 on page 11 of the revised manuscript.\nFigure S34. Free energy diagram for urea production on Ag(111)-Cu(111) with varying coupling and hydrogenation conditions. The insert shows an enlarged view of the pink region depicting the *CO +*NO2 →*CONOH transition.\nFigure S46. Free energy diagram of NO3RR(a) and HER(b) on Ag(111)-Cu(111).\nFigure S49. Free energy diagram of CO2-to-C2+ on Ag(111)-Cu(111).\nFigure S47. Structure diagrams of NO3RR, including the side view(a) and top view (b) on Ag(111)-Cu(111) surface. The blue, red, yellow, green and white balls represent Cu, Ag, O, N, H atoms, respectively.\nFigure S48. Structure diagrams of HER, including the side view and top view on Ag(111)-Cu(111) surface. The blue, red, and white balls represent Cu, Ag, H atoms, respectively.\nFigure S50. Structure diagrams of CO2-to-C2+, including the side view (a) and top view (b) on Ag(111)-Cu(111) surface. The blue, red, yellow, gray and white balls represent Cu, Ag, O, C and H atoms, respectively.\nTable S9. The free energy of the N-species intermediates involved in NO3RR on Ag(111)-Cu(111).\nStep System Free Energy (eV)\n2 *NO3 0.6375\n3 *NO3H 0.8735\n4 *NO2 -1.8368\n5 *NO2H -1.7767\n6 *NO -3.7938\n*NOH -4.0400\n7 *NHO -3.9986\n8 *NHOH -5.0556\n9 *NH -7.1555\n10 *NH2 -8.3352\n11 *NH3 -9.3675\nTable S10. The free energy of the intermediates involved in HER on Ag(111)-Cu(111).\nStep System Free Energy (eV)\n1 2 (H++e-) 0\n2 *H + (H++e-) -0.4641\n3 *H + *H -0.9366\nTable S11. The free energy of the C-species intermediates involved in CO2-to-C2+ on Ag(111)-Cu(111).\nStep System Free Energy (eV)\n2 *CO2 + *CO2 1.2647\n3 *COOH + *CO2 0.9801\n4 *CO + *CO2 -0.2371\n*CHO + *CO2 0.1391\n*CO + *COOH -0.5807\n5 *COCO -0.1873\n*COH + *CO2 0.5679\n*CHO + *COOH -0.1691\n*CO + *CO2 -1.6392\n6 *COCOOH -0.2813\n*COH + *COOH 0.1763\n*CHO + *CO2 -1.3560\n7 *COH + *CO2 -0.8370\n*COCO -0.7521\n*CHO + *CHO -0.9576\n*CH2O + *CO2 -1.9397\n8 *CHOH + *CO2 -1.5649\n*CHOCO -1.4261\n*CHO + *COH -0.5367\n*CH2O + *CHO -1.5215\n*CH2OH + *CO2 -2.3992\n9 *CH2OCO -0.7057\n*CH2O + *COH -1.1365\n*CH2OH + *CHO -2.2232\n10 *CH2 + *CO2 -3.1653\n*CH2OH + *COH -1.4841\n*CH2 + *CHO -2.8862\n*CH3 + *CO2 -4.5531\n11 *CH2CO -3.5253\n*CH2 + *COH -2.2778\n*CH2CHO -4.4546\n12 *CH2COH -3.9548\n*CH2CO -4.4641\n*CH4 + *CO2 -5.9225\n2) C-N coupling is a central challenge in urea synthesis, as it is crucial to the kinetics of the reaction. However, the study does not delve deeply enough into this issue. As mentioned in the paper, different nitrogen-containing intermediates, such as NO and NOH*, are generated during the nitrate reduction process, but the author only considers the coupling between NOH* and CO. While this coupling may indeed be experimentally observed, to highlight the superiority of this coupling, the interactions between other nitrogen-containing intermediates and CO or CO2 should also be considered. More importantly, the author mentions that this coupling can occur spontaneously, which may be inaccurate, as this explanation is solely based on thermodynamics. The author has not calculated the kinetic barriers for these reactions, and therefore, this aspect should also be investigated.\nAuthor response: We are profoundly grateful to the reviewer for their invaluable and perceptive suggestions and for addressing the misrepresentations in our work. In response to the reviewer's insightful comments, we have conducted a comprehensive analysis of the coupling trends among all potential reactions in the urea synthesis pathway. As illustrated in Figure S34, we evaluated the urea synthesis process on Ag(111)-Cu(111) (Ag67Cu33), beginning with the adsorption of CO2 and NO3-. Subsequently, each step was systematically assessed for possible couplings between N- and C-intermediates, as well as their respective hydrogenation processes. Notably, our results indicated that the free energy barrier for hydrogenation remained lower than that for coupling, until the critical formation of *CONOH from *NOH + *CO, thereby establishing the first C-N bond, which was consistent with our SR-FTIR findings. Furthermore, our analysis extended to subsequent hydrogenation steps, ultimately leading to the formation of the second C-N bond between *CONH2 and *NO2, resulting in the intermediate *H2NCONO2.\nSubsequently, we conducted a detailed investigation into the kinetic barriers associated with two critical C-N coupling reactions to elucidate their propensity to occur from a kinetic perspective. As illustrated in Figure S35, the first C-N coupling step involving *NOH and *CO exhibited a transition state (TS) energy barrier of 0.67 eV, notably lower than the 0.95 eV barrier for the protonation of *NOH. Thereafter, *CONH2 and *NO2 overcame a relatively modest TS energy barrier (0.87 eV), forming the intermediate *H2NCONO2. In contrast, the TS energy barrier for the transformation of *CONH2 + *NO2 into *CONH2 + *NOH was significantly higher (1.37 eV).\nConsidering the free energy diagram for urea production (Figure S34) and the kinetic barriers for the two C-N couplings (Figure S35) on Ag67Cu33, these C-N couplings were not only thermodynamically spontaneous but also kinetically feasible, rendering them more accessible compared to direct protonation (Adv. Mater. 2024, 36, 2402160; Nat. Commun. 2022, 13, 5337). We have thoroughly supplemented these details as well as corrected inappropriate expression in our revised manuscript (please refer to line 351 on page 11 of the revised manuscript).\nFigure S34. Free energy diagram for urea production on Ag(111)-Cu(111) with varying coupling and hydrogenation conditions. The insert shows an enlarged view of the pink region depicting the *CO +*NO2 →*CONOH transition.\nFigure S35. The reaction pathways and their corresponding potential-time of the system during the slow-growth process under constant potential method: (a-b) the first C-N coupling (forming *CONOH), (c-d) the protonation of *NOH (yielding *CO + *NHOH), (e-f) the second C-N coupling (forming *NOCONH2), and (g-h) the protonation of *NO (yielding *NOH + *CONH2). The insert was the structures corresponding to the initial, transition, and final states are illustrated. The blue, red, yellow, green, gray and white balls represent Cu, Ag, O, N, C, H atoms, respectively.\nTable S4. The free energy of adsorbent intermediates during urea synthesis on Ag(111)-Cu(111).\nStep System Free Energy (eV)\n2 *CO2 + NO3- 0.4477\n3 *CO2 + *NO3 1.3343\n4 *CO2 + *NO3H 1.4121\n5 *CO2 + *NO2 -1.2469\n*CO2NO2 -0.9396\n6 *COOH + *NO2 -1.5334\n*CO2 + *NO2H -1.4953\n*CO2 + *NO2 -2.7849\n7 *COOH + NO2H -1.4527\n*NO2COOH -0.7885\n*CHO + *NO2 -2.3436\n*CO2 + *NO2H -2.6933\n8 *COH + *NO2 -1.9606\n*CONO2 -1.6959\n*CHO + *NO2H -2.2508\n*CO2 + *NO2 -4.6828\n9 *COH + *NO2H -1.8802\n*CONO2H -2.2126\n*CHO + *NO2 -4.2631\n*CO2 + *NHO -4.7222\n10 *CO2 + *NOH -4.8814\n*COH + *NO2 -3.8627\n*CONO2 -4.3541\n*CO2 + *NHOH -5.0504\n11 *CONOH -5.5087\n12 *CON -8.4376\n13 *CONH -8.5656\n14 *CONH2 -9.5578\n15 *CONH2 + *NO3 -8.7556\n16 *CONH2 + *NO2 -11.2472\n17 *CONH2 + *NO2 -13.2716\n*H2NCONO2 -13.6563\n18 *CONH2 + *NOH -13.3963\n*CONH2 + *NHO -13.3089\n19 *H2NCONOH -13.7966\n20 *H2NCON -16.5031\n21 *H2NCONH -17.7077\n22 *H2NCONH2 -18.6478\n3) The author claims that the formation of AgCu heterojunctions facilitates the reduction of CO2 to CO. However, theoretical calculations show that the Gibbs free energy for this process is as high as 1.27 eV, which is higher than the 0.99 eV for traditional Cu catalysts. Based on this data, the AgCu catalyst does not exhibit good performance for CO formation, as it hinders the formation of CO*, which is detrimental to urea synthesis. This contradicts the experimental results. I suspect this may be related to the CHE model used in the calculations, as this model does not account for the effects of pH and potential, and the author also neglects the influence of solvent effects.\nAuthor response: We are profoundly indebted to the reviewer for their professional and invaluable insights concerning the optimization of our calculation method, which has significantly improved its reliability. In response to the reviewer's suggestions, we have employed the CP-VASP code (J. Am. Chem. Soc. 2021, 143, 9423-9428; J. Am. Chem. Soc. 2023, 145, 25352-25356) to conduct a thorough investigation into the influence of potential and pH on the C-N coupling process, while implicit solvation was effectively addressed through VASPsol++ (J. Chem. Phys. 2023, 159). As illustrated in Figure 4g, the revised free-energy diagram for CO2 reduction reaction (CO2RR) on Ag, Cu, and Ag67Cu33 catalysts demonstrated that the free energy barrier for converting CO2 into *CO was distinctly lowest on Ag(111)-Cu(111) (Ag67Cu33) with only 0.44 eV, which facilitated *CO formation for further C-N coupling, and notably lower than that of transitional Cu catalysts (0.99 eV). Moreover, the desorption of *CO from Ag67Cu33 presented significant challenges, requiring an energy input of approximately 0.46 eV, which was advantageous for stabilizing the *CO intermediate. Therefore, these effects collectively enhanced the surface coverage of *CO on Ag67Cu33, thereby substantially improving the catalytic efficiency for C-N coupling. Please refer to line 679 on page 18 of the revised manuscript for more details of calculation method.\nFigure 4g. The Free energy diagram of CO2RR on Ag(111), Ag(111)-Cu(111) and Cu(111).\nFigure S45. Structure diagrams of CO2RR. Top view and side view on Ag(111)-Cu(111)(a), Cu(111)(b), and Ag(111)(c). The blue, red, yellow, gray and white balls represent Cu, Ag, O, C, H atoms, respectively.\nTable S7. The free energy of adsorbed C-species intermediates in CO2RR.\nAg(111)-Cu(111) Cu(111) Ag(111)\nSystem Free Energy (eV) Free Energy (eV) Free Energy (eV)\nCO2 0 0 0\n*CO 0.4477 0.4936 0.5819\n*COOH 0.3701 0.0151 0.3445\n*CO -0.7127 -0.5518 0.0076\nCO -0.2557 -0.2557 -0.2557\n4) The author considers the impact of pH on catalytic activity in the computational method, with a selected pH value of 7. However, the CHE model oversimplifies the influence of pH on catalytic activity, often yielding results that contradict experimental observations. For detailed discussion, please refer to a series of articles by Zhiyao Duan et al. on calculating under constant potential. Therefore, to better account for the effect of pH, we suggest re-evaluating the results using the constant potential method.\nAuthor response: We would like to express our sincere gratitude for the insightful recommendations provided by the reviewer. Firstly, we utilized a pH meter to determine the pH value during the electrochemical experiments. Our results (Figure R1) indicated that the pH of the CO2-saturated electrolyte was consistently maintained at approximately 7.4 throughout the duration of testing, attributed to continuous CO2 injection. Subsequently, we employed a constant potential method (as previously detailed, please refer to line 679 on page 18 of the revised manuscript for more details of calculation method), establishing both pH and potential at 7.4 and -0.96 V vs. SHE (equivalent to -0.52 V vs. RHE). This methodology ensured alignment with the experimental conditions outlined in our study. Furthermore, we diligently undertook a comprehensive recalibration of all findings from our work (Figure S33-50 and Table S4-11), adequately considering the effects of pH and potential, thereby enhancing the reliability of our derived outcomes.\nFigure R1. The pH values of post-electrolytes at various potentials ranging from -0.12 to -0.62 V (vs. RHE) in an H-cell.\nFigure 4. Operando Raman and SR-FTIR measurements along with DFT calculations for urea synthesis mechanism. (d) Bader charge for Ag(111)-Cu(111). The red and blue colors indicated charge accumulation and depletion, respectively. (e) Free energy diagram of urea production on Ag(111)-Cu(111). (f) Charge density distributions of adsorbed *NOH on Ag(111)-Cu(111). The cyan and purple areas display the electron gain and loss, respectively (Isosurface= 0.005 e/Å3). The blue, red, yellow, green and white balls represent Cu, Ag, O, N, H atoms, respectively. Free energy diagram of CO2RR(g) and projected density of states (PDOS) of d orbitals(h) on Ag(111), Ag(111)-Cu(111) and Cu(111).\nFigure S38. Free energy diagram of urea synthesis on Cu(111).\nFigure S41. Free energy diagram of urea synthesis on Ag(111).\nTable S5. The free energy of adsorbent intermediates during urea synthesis on Cu(111).\nStep System Free Energy (eV)\n2 *CO2 + NO3- 0.4936\n3 *CO2 + *NO3 1.0856\n4 *CO2 + *NO3H 1.3163\n5 *CO2 + *NO2 -1.4726\n6 *COOH + *NO2 -1.8759\n7 *CO2 + *NO2 -2.6141\n8 *CO2 + *NO2H -2.4600\n9 *CO2 + *NO2 -4.6495\n*CONO2 -4.5449\n10 *CO2 + *NHO -5.2076\n*CO2 + *NOH -5.1284\n*CO2 + *NHOH -5.8470\n11 *CONHO -5.5440\n*CO2 + *NH -8.2306\n12 *CONHOH -6.0989\n*CO2 + *NH2 -9.1397\n13 *CONH -8.4934\n*CO2 + *NH3 -10.0871\n14 *CONH2 -9.5056\nTable S6. The free energy of adsorbent intermediates during urea synthesis on Ag(111).\nStep System Free Energy (eV)\n2 *CO2 + NO3- 0.5818\n3 *CO2 + *NO3 0.9984\n4 *CO2 + *NO3H 1.3241\n5 *CO2 + *NO2 -1.4073\n6 *COOH + *NO2 -1.6837\n7 *CO2 + *NO2 -2.0732\n8 *CO2 + *NO2H -1.9172\n9 *CO2 + *NO2 -3.8233\n*CONO2 -4.1147\n10 *CO2 + *NOH -3.9534\n*CO2 + *NHO -4.2439\n*CONHO -5.2563\n11 *CO2 + *NHOH -4.8808\n12 *CONHOH -5.7213\n13 *CONH -8.2742\n14 *CONH2 -9.1902\n15 *CONH2 + *NO3 -8.7604\n16 *CONH2 + *NO2 -11.1708\n17 *CONH2 + *NO2 -12.6162\n18 *H2NCONO2 -13.6909\n19 *H2NCONOH -14.5160\n20 *H2NCON -16.4564\n21 *H2NCONH -17.5753\n22 *H2NCONH2 -18.3788\n5) To make readers better understand the AgCu structure, the corresponding contcar files computed by VASP should be provided, including catalyst, adsorbed intermediates.\nAuthor response: We would like to express our sincere gratitude to the reviewer for their thorough and insightful feedback. To enhance readers' understanding of the AgCu structure and the details of our calculations, we have incorporated both top and side views of the models produced during the DFT calculation processes in the revised supplementary information (Figure S33-50). In accordance with the data policy of Nature Communications, we have provided computational models including detailed CONTCAR files for catalysts and adsorbed intermediates generated using VASP in https://doi.org/10.6084/m9.figshare.28636625.v2 (this link was in Data availability on page 23).\nReviewer #3: In this article, Feng et al. constructed analogous rectifying interfaces in AgCu biphasic aerogels with stable Cuδ+ active sites, which simultaneously promoted *NOH adsorption and *CO coverage, thereby improving C-N coupling dynamics. Specifically, using Ag67Cu33 in a flow cell, ultrahigh urea yield (54.8 mmol h-1 gcat.-1) with notable Faradaic efficiency (FE, 36.6%) at a low potential (-0.52 V vs. RHE) was achieved. Overall the manuscript is quite impressive. Besides, the manuscript can be revised to address some issues before being considered for publication in this journal.\nAuthor response: We are deeply grateful to the reviewer for dedicating their valuable time to thoroughly review our work and offering insightful suggestions, which have greatly enriched our manuscript, and your acknowledgment of our efforts is encouraging. We have carefully examined your comments and conducted additional experiments and analyses to address each point raised by the reviewer in detail. Our ultimate objective is to improve the overall quality of our manuscript through the addition of robust experimental evidence and detailed explanations.\nIn response to the reviewer's constructive feedback, we have incorporated the following revisions:\n1. Reaction conditions: Through a series of experiments, we systematically explored the impact of varying reaction parameters, including NaBH4 concentrations, reaction time and temperature, on the structure, morphology and catalytic performance of the synthesized catalysts.\n2. Potential range: We have conducted potentiostatic tests on Ag67Cu33 in an H-cell, spanning a potential range from -0.12 V to -0.72 V (vs. RHE), and subsequently analyzed all products across the different potentials.\n3. Control experiments: To further corroborate the critical role of the Ag-Cu interfaces in enhancing catalytic efficiency, we have meticulously synthesized pure Ag, pure Cu and physically mixed Ag2Cu1 catalysts, and conducted control experiments to assess their urea synthesis performance.\n4. XANES characterization: We examined the chemical states of Cu in the Ag67Cu33 catalyst through synchrotron radiation X-ray absorption near-edge structure analysis, confirming that Cu retained its metallic state during the electrochemical process.\n5. Catalyst leaching investigation: We re-evaluated the stability in both three-electrode and two-electrode flow cells, employing ICP-MS to analyze the concentrations of Ag and Cu in the post-electrolytes as well as the catalysts on carbon paper after the tests.\n6. Discussion of state-of-the-art catalysts: We have thoroughly reviewed the latest advancements in electrocatalytic urea synthesis, and expanded our discussion to include metal-N4 and metal phthalocyanine catalysts in the revised manuscript.\n7. pH value analysis: Utilizing a precision pH meter, we monitored and recorded the variations in pH values of the post-electrolytes after potentiostatic tests in an H-cell, as well as stability tests conducted in both three-electrode and two-electrode flow cells.\n8. UV data: We have comprehensively integrated all UV data corresponding to the results presented in the manuscript, now provided in the Supporting Information.\n9. Influence of the crystallographic alignment: We constructed models encompassing diverse Cu crystallographic orientations and conducted DFT calculations on key reaction steps, thereby elucidating the influence of distinct Cu facets on urea synthesis in conjunction with insights from previous studies.\n10. KNO3 concentration selection: We performed a series of experiments utilizing varying concentrations of KNO3 in a three-electrode flow cell on Ag67Cu33 catalyst, subsequently evaluating their catalytic performance for urea synthesis.\nWe once again extend our deepest gratitude to the esteemed reviewer for your invaluable and insightful guidance. We sincerely hope that our detailed explanations and the thoroughly revised manuscript will fully meet your expectations. We look forward to receiving your further feedback.\nThe following sections provide thorough responses to each of reviewer's inestimable recommendations.\n1. How do the specific concentrations of NaBH4 and the reaction conditions influence the formation of the AgCu biphasic aerogels? Could variations in these parameters affect reproducibility or performance consistency?\nAuthor response: We are profoundly grateful for the invaluable and insightful feedback from the reviewer. Based on prior researches on metallic aerogels, we successfully synthesized AgCu aerogels via a NaBH4 reduction process (with a molar ratio of NaBH4 to metal ions set at R/M = 5:1) conducted in an ice-water bath. After approximately 35 minutes, this method yielded black precipitates. The carefully chosen reaction conditions were designed not only to control the reaction rate and ensure the reaction proceeded smoothly, but also toguarantee the complete reduction of metal ions. In accordance with the reviewer's recommendations, we conducted a series of meticulously designed experiments utilizing controlled variables to investigate the effects of reaction conditions, including NaBH4 concentrations, reaction temperature and reaction time, on the structure, morphology and catalytic performance of the synthesized catalysts. To ensure consistency and minimize potential human error, we standardized the preparation of metal ion solutions in all subsequent synthesis experiments, and selected Ag67Cu33 as a representative sample, following a theoretical molar ratio of Ag:Cu at 67:33.\nInitially, we conducted an in-depth investigation into the impact of varying NaBH4 concentrations. We systematically adjusted the molar ratio of NaBH4 to metal ions (R/M) across a spectrum from 1:1 to 20:1,specifically including ratios of 1:1, 2:1, 5:1, 10:1, and 20:1. Notably, NaBH4 underwent hydrolysis in neutral water, with its rate significantly accelerating as temperature rose. The detailed reactions proceed as follows: NaBH4 → BH4- + Na+, BH4- +2H2O → BO2- +4H2, NaBO2 + H2O → NaOH + H3BO3, which rendered the solution mildly basic and could be catalyzed by metal ions (Ag+ and Cu2+) (J. Saudi Chem. Soc. 2021, 25, 101258). The XRD analysis of catalysts at different R/M ratios, illustrated in Figure R2, revealed that catalysts at R/M ratios of 5:1 and 10:1 exhibited superior crystallinity without any impurities such as Cu2O or Cu(OH)2. SEM images (Figure R3) further confirmed that all catalysts possessed a characteristic three-dimensional reticular aerogel structure, with increasing diameters observed as NaBH4 concentration increased. ICP-MS results (Table R1) indicated a notably lower Cu content at an R/M ratio of 1:1, while a slightly higher Cu content was detected at 20:1. Combining these observations with experimental phenomena, we attributed the observed differences to several possible factors (J. Chem. Educ. 2007, 84, 322; Carbohydr. Polym. 2014, 107, 167-173): when R/M ≤ 2, insufficient BH4- were available to stabilize the reduced metal, leading to rapid aggregation and sedimentation within 15 minutes, resulting in incomplete reduction in the mildly alkaline environment; when 2 < R/M ≤ 10, BH4- effectively facilitated the formation of colloidal metal particles, controlling the aggregation rate and promoting the development of larger diameters and complete reduction; at an R/M ratio of 20:1, the reaction proceeded extremely rapidly within 3 minutes due to the excessive NaBH4, which increased the overall ionic strength and led to aggregation, potentially resulting in partially segregated Ag and Cu (resembling a physical mixture) due to numerous reaction centers with fast rate.\nFigure R2. The XRD patterns of catalysts with various R/M: 1:1, 2:1, 5:1, 10:1, 20:1.\nFigure R3. The SEM images of catalysts with various R/M: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e).\nTable R1. The ICP-MS results of catalysts with various R/M: 1:1, 2:1, 5:1, 10:1, 20:1.\nNaBH4 concentration (Molar ratio) 1:1 2:1 5:1 10:1 20:1\nPractical Ag:Cu 78.0:22.0 63.1:36.9 66.6:33.4 65.3:34.7 61.9:38.1\nSubsequently, we systematically controlled the reaction conditions by separately adjusting the temperature (ice-water bath at 0 ℃, ambient temperature at 25 ℃, and elevated temperature at 50 ℃) and the reaction time (1 h, 2 h, 4 h). As illustrated in Figure R4, the XRD patterns exhibited minimal differences, with the notable exception being the formation of CuO when the reaction was conducted at 50 ℃. SEM images (Figure R5) revealed that increasing the temperature led to a significant enlargement of catalyst's diameter, culminating in agglomerated structures that compromised the aerogel's integrity at 50 ℃ (Sci. Rep-UK 2019, 9, 15212). In contrast, extending the reaction time did not result in substantial changes, apart from minor aggregations. ICP analysis (Table R2) indicated that the molar ratio of Ag to Cu remained remarkably consistent across different reaction times, while slight variations were observed under different temperature conditions. In summary, prolonging the reaction time had limited impact and potentially induced undesirable aggregations, whereas elevating the temperature accelerated the reaction rate and the hydrolysis of NaBH4, leading to larger particle diameters and other adverse side reactions.\nFigure R4. The XRD patterns of catalysts under different reaction temperature(a) and reaction time(b).\nFigure R5. The SEM images of catalysts under different reaction temperature: 0 ℃/1 h(a), 25 ℃/1 h(b), 50 ℃/1 h(c) and reaction time: 0 ℃/2 h(d), 0 ℃/4 h(e).\nTable R2. The ICP-MS results of catalysts under different reaction temperature and reaction time.\nReaction conditions 0 ℃ / 1 h 0 ℃ / 2 h 0 ℃ / 4 h 25 ℃ / 1 h 50 ℃ / 1 h\nPractical Ag:Cu 66.6:33.4 66.9:33.1 66.8:33.2 61.2:38.8 64.1:35.9\nSubsequently, we conducted a series of experiments using an H-cell to rigorously evaluate the catalytic performance of these catalysts in urea synthesis, following a standardized protocol. As depicted in Figure R6, the sample with R/M = 5:1 exhibited superior urea yield rate and Faradaic efficiency compared to others. Considering the electrochemical surface area (ECSA) as indicated by the Cdl values (Figure R11) and the current density (Figure R8), it became apparent that the R/M = 5:1 sample possessed enhanced intrinsic catalytic activity despite a slightly lower current density. Combined with our previous research, this observation indicated that varying concentrations of NaBH4 might significantly influence the formation of Ag-Cu interfaces, which were critical for inducing Cuδ+ active sites for urea synthesis. Furthermore, post-test XRD analysis (Figure R13) revealed the absence of oxides or hydroxides across all samples, attributed to the LSV activation prior to testing. SEM images (Figure R14) confirmed the preservation of the aerogel structure. However, samples with R/M ≤ 2 exhibited increased diameters due to their tendency to agglomerate during catalytic reactions, likely driven by their high specific surface energy.\nFigure R6. The Urea yield rate(a) and Faradaic efficiency(b) of catalysts with various R/M.\nAs the reaction temperature increased, the performance of urea synthesis deteriorated, particularly at 50 ℃ (Figure R7). Although XRD analysis (Figure R15) confirmed that the overall structure remained unchanged after testing when compared to other samples, SEM observations (Figure R15) revealed that particle agglomeration resulted in an increased particle diameter and partial degradation of the aerogel structure at this elevated temperature. ECSA measurements (Figure R12) indicated a decrease in the Cdl value with rising temperature. In contrast, extended reaction time had little impact on the catalyst's structure, morphology, and catalytic efficiency, as evidenced by consistent results in urea yield, FE, ECSA, XRD, and SEM analyses (Figures R7, 12 and 16). However, under alkaline conditions, it was advisable to limit the duration of the reaction to avoid potential adverse effects (Carbohydr. Polym. 2014, 107, 167-173).\nIn summary, varying concentrations of NaBH4 significantly influenced the aerogel formation process, which in turn exerted a profound influence on the development of Ag-Cu interfaces, thus resulting in their distinct catalytic behaviors. Additionally, an increase in reaction temperature accelerated the reaction rate, leading to larger particle diameters and potential degradation of the aerogel structure, thereby compromising its catalytic efficiency. The aerogel synthesis was completed within 30 to 40 minutes, so extending the reaction time beyond this period offered minimal benefits, particularly in weakly alkaline environments, where excessive duration was neither advisable nor necessary. Consequently, our investigation had determined the optimal reaction conditions as follows: a ratio of 2 < R/M ≤ 10, a temperature range of 0 to 25 ℃, and a reaction time of 1 to 2 h. Under these conditions, the catalyst's structure, morphology, and catalytic performance were maintained without obvious alteration. Among these, a ratio of R/M = 5:1, conducted in an ice-water bath for 1 hour, was identified as the optimal condition in our work.\nFigure R7. The Urea yield rate(a) and Faradaic efficiency(b) of catalysts under different reaction temperature and reaction time.\nFigure R8. The i-t curves of catalysts with various R/M: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e).\nFigure R9. The i-t curves of catalysts under different reaction temperature: 0 ℃/1 h(a), 25 ℃/1 h(b), 50 ℃/1 h(c) and reaction time: 0 ℃/2 h(d), 0 ℃/4 h(e).\nFigure R10. The UV data of catalysts with various R/M: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e), different reaction temperature: 0 ℃/1 h(c), 25 ℃/1 h(f), 50 ℃/1 h(g) and reaction time: 0 ℃/2 h(h), 0 ℃/4 h(i).\nFigure R11. CV curves of catalysts with various R/M: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e), at the scan rate from 40 to 140 mV s-1. (f) Δj/2 plotted against the scan rates over different samples, where the slope is Cdl.\nFigure R12. CV curves of catalysts under different reaction temperature: 0 ℃/1 h(a), 25 ℃/1 h(b), 50 ℃/1 h(c) and reaction time: 0 ℃/2 h(d), 0 ℃/4 h(e), at the scan rate from 40 to 140 mV s-1. (f) Δj/2 plotted against the scan rates over different samples, where the slope is Cdl.\nFigure R13 The XRD patterns of catalysts with various R/M after test: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e).\nFigure R14. The SEM images of catalysts with various R/M after test: 1:1(a), 2:1(b), 5:1(c), 10:1(d), 20:1(e).\nFigure R15. The XRD patterns and SEM images of catalysts under different reaction temperature: 0 ℃/1 h(a, d), 25 ℃/1 h(b, e), 50 ℃/1 h(c, f).\nFigure R16. The XRD patterns and SEM images of catalysts under different reaction time: 0 ℃/1 h(a, d), 0 ℃/2 h(b, e), 0 ℃/4 h(c, f).\n2. The LSV figure shows that the reduction starts at nearly -0.1V. Why were specific potential ranges (-0.22 V to -0.62 V vs. RHE) selected for electrocatalysis? Could alternative ranges provide insights into different catalytic behaviours?\nAuthor response: We are deeply appreciative of your feedback regarding the selection of potential intervals, which was a critical aspect in electrochemical testing. Initially, we evaluated the catalysts within an H-cell by applying potentials ranging from -0.22 to -0.72 V (vs. RHE). As shown in Figure R17, the urea yield rate exhibited a volcano-type trend, with the optimal catalyst (Ag67Cu33) achieving peak performance at -0.52 V. Regarding Faradaic efficiency (FE), minimal variation was observed at lower potentials, while beyond -0.52 V, FE showed a downward trend. Corroborated by by-product analysis (Figure S9), this decline might be attributed to the enhanced competitive NO3RR at more negative potentials (Nat. Synth. 2023, 2, 624), thereby diminishing urea FE. Consequently, we narrowed our focus to the interval between -0.22 and -0.62 V (vs. RHE) for urea electrocatalysis.\nIn response to the reviewer's suggestion that reduction commenced near -0.1 V, we extended our investigation on Ag67Cu33 under applied potentials ranging from -0.12 to -0.72 V (vs. RHE) within an H-cell. As illustrated in Figure R18, the production of urea was confirmed even at a potential as low as -0.12 V.However, the urea yield rate was constrained by the relatively low current density. As the potential became increasingly negative, there was a significant escalation in the urea yield rate, along with a moderate increase in FE. The peak urea yield rate (25.7 mmol h-1 gcat.-1) and maximum FE (30.6%) were both achieved at -0.52 V. Subsequently, despite continued increases in current densities, the urea yield rate experienced a slight downturn, paralleled by a marked decrease in FE. In conjunction with the FE of all products (Figure R18c), the catalytic behavior of Ag67Cu33 remained relatively stable at lower potentials. However, at more negative potentials, the increased presence of N-bearing by-products indicated that the accelerated competitive NO3RR likely occupied more active sites (Chem. Eng. J. 2024, 485, 150052; Angew. Chem. Int. Ed.63, e202406515), which resulted in a reduced FE of urea at higher potentials.\nFigure R17. Urea yield rate(a) and corresponding FE(b) on AgxCu100-x at various applied potentials ranging from -0.22 to -0.72 V (vs. RHE).\nFigure R18. I-t curves(a), urea yield rate with corresponding FE(b), FE of all products(c), and the corresponding UV data for urea(d), NH3(e) and NO2-(f) in post-electrolytes obtained from Ag67Cu33 at various applied potentials ranging from -0.12 to -0.72 V (vs. RHE).\n3. Have the authors considered conducting control experiments using pure Ag and pure Cu catalysts in physically mixed configurations instead of biphasic aerogels to further validate the role of the affinity-modifying interface? Also, the authors should provide control experiments for urea synthesis using pure Ag and pure Cu catalysts.\nAuthor response: We are profoundly appreciative of your invaluable suggestions regarding the control experiments that underscore the pivotal role of Ag-Cu interfaces. Initially, we synthesized pure Ag and pure Cu aerogel catalysts using an identical methodology, with the only difference being the addition of either AgNO3 or Cu(NO3)2. Subsequently, we prepared a uniform physical mixture at a molar ratio of 2:1 to synthesize the Ag2Cu1 catalyst devoid of Ag-Cu interfaces (more details in Materials and Methods on page 14 of the revised manuscript). As shown in Figure S52, its urea synthesis performance was significantly inferior to that of Ag67Cu33, with no notable enhancement observed, thereby emphasizing the essential importance of the analogous rectifying interface effect within Ag-Cu biphasic aerogels. Moreover, we conducted additional experiments on pure Ag and pure Cu aerogel catalysts. As depicted in Figures 2c and 2d, all AgCu biphasic catalysts demonstrated markedly superior urea performance compared to pure Ag and pure Cu catalysts, thus confirming the indispensable role of the analogous rectifying interface effect in enhancing urea synthesis. Furthermore, the subpar urea yield rate from pure Ag suggested that the performance enhancement was not attributed solely to the introduction of Ag. Additionally, the Faradaic efficiency (FE) of all potential products (Figure R19) revealed higher FE for NO2- and NH3 on pure Cu, indicating that increased competition from NO3RR diminished urea FE, which aligned with DFT calculation results (Figure S38), suggesting that the analogous rectifying interface effect might mitigate competitive reactions. Please refer to line 188 on page 5 and line 386 on page 11 of the revised manuscript, where further details are provided.\nFigure R19. Faradaic efficiencies for all products on Ag(a), Ag80Cu20 (b), Ag67Cu33 (c), Ag50Cu50 (d), Cu(e).\nFigure S52. LSV curves(a), i-t curves(b), Faradaic efficiencies for all products(c) on Ag2Cu1 at different applied potentials. (d) Comparison of the urea yield rate of Ag2Cu1 with other catalysts. The corresponding UV spectroscopy data for urea(e), NH3(f) and NO2-(g) on Ag2Cu1.\n4. Additional surface or bulk characterization techniques can be considered (e.g., XANES, EXAFS) to further confirm the oxidation state of Cu and Ag during the electrocatalytic process.\nAuthor response: We sincerely appreciate the insightful recommendations provided by the reviewer concerning supplementary characterization to elucidate the chemical states in AgCu catalysts. In regards to the XPS analysis results of Ag67Cu33 (Figure 1g-h), it was evident that Ag was present in its elemental form with a zero-valent state, while Cu showed a minor oxidation due to exposure to air. Therefore, to investigate the chemical state of Cu during potentiostatic tests, we performed additional ex-situ synchrotron radiation X-ray absorption near-edge structure (XANES) analysis of Ag67Cu33 at the National Synchrotron Radiation Laboratory (NSRL), and we would like to acknowledge the support provided by the BL10B in NSRL for the characterization experiments using Synchrotron Radiation.\nWe evaluated the Ag67Cu33 catalyst following LSV activation and potentiostatic tests (from -0.22 V to -0.62 V). Post-testing, the catalyst was promptly transferred to a 25°C Ar atmosphere for drying, and subsequently characterized using XANES to probe the Cu L-edge (Figure S30a), thereby minimizing potential air oxidation. As illustrated in Figure S30b, the Cu L-edge spectrum recorded at open-circuit potential (OCP) exhibited distinct peaks at 931.9 eV and 951.8 eV, attributed to the 2p3/2 and 2p1/2 states of metallic Cu, while a minor peak indicative of Cu-O coordination was also observed (Nat. Commun. 2024, 15, 2816; Nat. Catal. 2018, 1, 421-428), likely arising from surface oxidation during the catalyst loading process on carbon paper. Notably, following subsequent electrochemical treatment, the spectral features only exhibited metallic Cu peaks, without oxide signatures. Furthermore, the integration of operando Raman spectroscopy data (Figure S31), XPS results, and Cu L-edge XANES spectra confirmed the preservation of metallic Ag and Cu in Ag67Cu33 during the electrochemical processes. Please refer to line 283 on page 8 of the revised manuscript.\nFigure S30. (a) The set-up of XANES measurement system. (b) The Cu L-edge XANES spectra of Ag67Cu33 after LSV activation and potentiostatic tests conducted from -0.22 V to -0.62 V.\n5. Stability tests show no significant morphological changes, but what about the possible leaching of Ag or Cu? ICP-MS can be used to confirm the absence of dissolved metals for long-term use.\nAuthor response: We are profoundly appreciative of the reviewer for highlighting the potential leaching of the catalyst, which may influence the stability tests. Following the reviewer's insightful suggestions, we conducted additional stability assessments in both three-electrode and two-electrode flow cells (Figure R20). We employed ICP-MS to rigorously analyze the concentrations of Ag and Cu in the post-electrolytes and the catalysts deposited on carbon paper after tests. As shown in Figure S23, a trace amount of Ag and Cu was detected in the post-electrolytes, with the final catalyst mass on carbon paper stabilizing at 0.97 mg (closely aligning with the calculated loss of 0.95 mg from the electrolytes). This indicated minimal loss during the stability tests, starting from an initial catalyst loading of 1 mg. Notably, the molar ratio of Ag to Cu in the post-test catalyst (71.9:28.1) was higher than the initial ratio of Ag67Cu33 (66.6:33.4), consistent with the greater loss of Cu observed in the post-electrolytes. Consequently, we inferred that the primary cause of catalyst loss was attributed to surface oxidation at the open circuit potential (OCP) when refreshing the electrolytes (Nat. Commun. 2024, 15, 10331). Then under reductive potentials, this surface oxide might dissolve, leading to catalyst loss, which was supported by in-situ Raman spectroscopy results (Figure S31). Similar observations were made at the cathode in the two-electrode flow cell (Figure S59), where the post-test catalyst on carbon paper had a loading of 0.44 mg (Ag:Cu = 74.4:25.6), with more Cu lost in the electrolytes (calculated retention of 0.43 mg). In contrast, the anode exhibited a molar ratio of Ag to Cu in the post-test catalyst on carbon paper (loading 0.43 mg, calculated retention of 0.42 mg) of 69.8:30.2, close to the initial ratio. This indicated that despite formaldehyde (HCHO) having a lower oxidation potential, some catalyst oxidation and dissolution still occurred under anodic conditions (Nat. Commun. 14, 525). Overall, despite a slight leaching of Ag and Cu observed during stability tests, the catalyst did not dissolve rapidly or extensively, and a significant portion remained intact post-testing, thereby preserving its performance throughout the testing period. We attributed this remarkable structural stability to two key factors (Matter 2021, 4, 54-94): firstly, the three-dimensional porous network structure inherent to metal aerogels, which exhibited intrinsic thermodynamic and chemical robustness, rendering it highly suitable for various catalytic applications; secondly, the incorporation of Ag as a secondary metal, owing to its stable chemical properties, significantly enhanced the material's oxidation resistance and mitigated leaching. Please refer to line 262 on page 7 and line 447 on page 13 of the revised manuscript, where further details are provided.\nFigure R20. The performance of urea synthesis during stability tests in three-electrode(a) and two-electrode(b-c) flow cells using Ag67Cu33, along with corresponding UV data for urea in three-electrode(d) and two-electrode(e) flow cells.\nFigure S23. The catalyst loading of Ag67Cu33 during stability tests in the three-electrode flow cell, with the inset table showing the Ag and Cu content in the post-electrolytes.\nFigure S59. The catalyst loading of Ag67Cu33 during stability tests in the two-electrode flow cell, with the inset table showing the Ag and Cu content in the post-electrolytes.\n6. The authors discuss electrochemical urea synthesis, a promising alternative to the traditional urea synthesis process using bimetallic systems. They should add some discussion from recent literature to keep the manuscript up-to-date using catalysts based on metal-N4 or metal phthalocyanine catalysts.\nAuthor response: We are deeply grateful to the reviewer for their invaluable suggestions, which have been instrumental in ensuring that our manuscript remains at the forefront of recent research. In response to the reviewer's insightful comments, we have incorporated the discussion on metal-N4 and metal phthalocyanine catalysts that exhibited high efficiency in urea synthesis (Angew. Chem. Int. Ed. 63, e202410625; Nat. Sustain. 7, 442-451; Nat. Commun. 15, 8858). Please refer to line 48 and line 62 on page 2 of the revised manuscript, where further details are provided.\n7. The pH value of the electrolyte may be varied accompanied with ammonia production (by-product from NO3– reduction). How the catalytic stability can be maintained?\nAuthor response: We express our gratitude for your insightful observations regarding the potential issues stemming from pH fluctuations during the electrolysis process. To comprehensively investigate these pH changes, we conducted a series of measurements on post-electrolyte pH values at various potentials ranging from -0.12 to -0.72 V (vs. RHE) in an H-cell over Ag67Cu33. As illustrated in Figure R21, the initial pH of the 0.5 M CO2-saturated KHCO3 solution (containing 2000 ppm KNO3) was accurately measured at 7.35. Upon applying increasingly negative potentials, the pH values exhibited a slight upward trend but remained below 8, which value was observed in a standard 0.5 M KHCO3 solution (with 2000 ppm KNO3). Notably, as the pH increased under more negative conditions, there was a corresponding rise in NH3 production (Figure R18). Thus, this phenomenon could be attributed to the *NO2 →*NH3 reaction, which rapidly consumed protons and was the principal driver behind the rapid pH alterations during NO3RR (Adv. Mater. 2024, 36, 2304021).\nIn subsequent experiments, we assessed the pH changes of post-electrolytes in both three-electrode and two-electrode flow cell (Figures R22), revealing the negligible pH variations throughout these stability tests. We attributed this remarkable stability to two key factors: firstly, the applied potential in stability tests was set at -0.52 V (vs. RHE), resulting in only a modest amount of NH3 being generated. More crucially, the continuous introduction of CO2 throughout the process could react with H2O to form carbonic acid (CO2 + H2O ⇌ H2CO3), which could subsequently decompose into bicarbonate ions and hydrogen ions (H2CO3 ⇌ HCO3- + H+). This dynamic equilibrium helped maintain the pH within the range observed in a 0.5 M KHCO3 solution (pH = 8), thereby ensuring that pH fluctuations did not compromise catalytic performance during electrochemical process.\nFigure R21. The pH values of post-electrolytes at various potentials ranging from -0.12 to -0.72 V (vs. RHE) in an H-cell.\nFigure R22. The pH values of post-electrolytes during stability tests in a three-electrode(a) and two-electrode(b) flow cell over Ag67Cu33.\n8. The authors carry out the DAMO method and Nessler's reagent test to detect urea and ammonia if produced from NO3– reduction. However, all the UV data for the samples are missing in the manuscript as well as in the supporting information. The author should add that missing information.\nAuthor response: We are deeply appreciative of the reviewer's insightful recommendation to enhance our work by incorporating additional UV data in the Supporting Information. In this study, following the dilution of the electrolyte to an optimal concentration, we utilized UV-Vis spectroscopy for a quantitative analysis of urea, NO2-, and NH3 production (more details in Product quantification of the revised manuscript). Heeding your valuable suggestion, we have now included all relevant UV data corresponding to the results presented in the manuscript within the Supporting Information. For your reference, please see the revised Supporting Information, where the comprehensive data can be found.\nFigure S8. Urea yield rate, Faradaic efficiency, and corresponding UV data for urea on Ag(a-b), Ag80Cu20 (c-d), Ag67Cu33 (e-f), Ag50Cu50 (g-h), Cu(i-j) under various applied potentials.\nFigure S9. Faradaic efficiencies for all products on Ag(a), Ag80Cu20 (e), Ag67Cu33 (i), Ag50Cu50 (m), Cu(q) at different applied potentials. The corresponding UV spectroscopy data for urea, NH3 and NO2- on Ag(b-d), Ag80Cu20 (f-h), Ag67Cu33 (j-l), Ag50Cu50 (n-p), Cu(r-t), respectively.\nFigure S10. The corresponding UV spectroscopy data for urea on Ag67Cu33 in stability tests at -0.52 V for 6 h.\nFigure S16. LSV curves(a), i-t curves(b), Faradaic efficiencies for all products(c), urea yield rate(d) at -0.52 (V vs. RHE), and corresponding UV spectroscopy data for urea(e), NH3 and NO2-(g) on Ag67Cu33 in different cathode electrolytes.\nFigure S17. The UV spectroscopy data(a) and 1H NMR spectra(b) corresponding to the quantification of urea by using the diacetyl monoxime method and 1H NMR method, respectively.\nFigure S18. (a) The actual test process of a three-electrode flow cell. The i-t curves(b), UV-vis absorption spectra for urea(c), together with the corresponding urea yield rate and Faradaic efficiency(d) of Ag67Cu33 at -0.52 V under varying KNO3 concentrations.\nFigure S19. I-t curves(a) and the corresponding UV data for urea in post-electrolytes(b) obtained on Ag67Cu33 at different applied potentials.\nFigure S20. The corresponding UV data for urea(a), NH3(b) and NO2-(c) in post-electrolytes obtained on Ag67Cu33.\nFigure S21. The corresponding UV spectroscopy data for urea on Ag67Cu33 in stability tests at -0.52 V for 24 h.\nFigure S52. LSV curves(a), i-t curves(b), Faradaic efficiencies for all products(c) on Ag2Cu1 at different applied potentials. (d) Comparison of the urea yield rate of Ag2Cu1 with other catalysts. The corresponding UV spectroscopy data for urea(e), NH3(f) and NO2-(g) on Ag2Cu1.\nFigure S53. The E-t curves(a) and the corresponding UV data for urea in post-electrolytes(b) obtained on Ag67Cu33 in the two-electrode system ranging from 15 mA cm-2 to 45 mA cm-2.\nFigure S55. The urea yield rate and corresponding FE in the cathode(a) and HCOOH yield rate and corresponding FE in the anode(b) of Ag67Cu33 in the two-electrode system at 40 mA cm-2 by 10 cycles, with each cycle presenting 1 h. (c) The corresponding UV data for urea in post-electrolytes obtained at cathode.\nFigure S57. The E-t curves(a) and corresponding UV data for urea in post-electrolytes obtained on Ag67Cu33 in the two-electrode system at 40 mA cm-2 of long-term electrolysis for 60 h.\n9. Density functional theory (DFT) models are based on specific crystallographic orientations. How would variations in the crystallographic alignment of the Ag and Cu interfaces affect the results?\nAuthor response: We extend our sincere gratitude to the reviewers for their thorough consideration about the crystallographic orientations influencing the results. Previous studies have demonstrated that different Cu facets exhibit varying catalytic activities in CO2 reduction reactions (J. Am. Chem. Soc. 2022, 144, 259-269), NO3- reduction reactions (Angew. Chem. Int. Ed. 2023, 62, e202303327) and urea synthesis (Nat. Commun. 2024, 15, 1095). In our work, we observed that Ag(111) and Cu(111) appeared as the strongest peak in XRD pattern, supported by observations from HR-TEM and HAADF-STEM (Figure 1), which were also the thermodynamically favorable facets (Nat. Commun. 2024, 15, 1749). As a result, we selected Ag(111) and Cu(111) as structural models for Density Functional Theory (DFT) calculations. According to prior investigations, the surface energy for Cu crystal facets (ACS Nano 2016, 10, 4559-4564) ranked from highest to lowest, was Cu(110), Cu(100), and Cu(111), while the high surface energy was likely associated with highly active sites (Joule 2020, 4, 2562-2598). However, predicting the behavior of Ag-Cu(hkl) with analogous rectifying interface effect on urea synthesis process, which involved complex 16 electron transfers, remained challenging. Therefore, to address the insightful suggestions from the reviewers more effectively, we developed models incorporating various Cu crystallographic orientations with Ag(111) to construct the Ag-Cu interface, based on the understanding that Cu sites served as the active centers in urea synthesis (Figure R23). Subsequently, we conducted DFT calculations on Ag(111)-Cu(100), Ag(111)-Cu(110), and Ag(111)-Cu(111) to systematically evaluate critical steps, including CO2 and NO3- adsorption, the first C-N coupling, and the second C-N coupling, aiming to investigate their influence on catalytic performance.\nThe computational models including detailed CONTCAR files for catalysts and adsorbed intermediates were in https://doi.org/10.6084/m9.figshare.28636625.v2 (in Data availability on page 23).\nAs shown in Figure R22 and Table R3, Ag(111)-Cu(100) and Ag(111)-Cu(110) exhibited significantly reduced adsorption energies for CO2 and NO3-. However, during the first C-N coupling reaction, Ag(111)-Cu(110) required an energy input of 1.54 eV. In contrast, for the Ag(111)-Cu(100) system, both the first and second C-N couplings showed the largest Gibbs free energy change values, suggesting a potential preference for urea synthesis. These results were consistent with Xu's work on DFT calculations regarding the transformation of NO3- + CO2 into urea (Nat. Commun. 2024, 15, 1095). According to previous studies, Cu(100) has been widely recognized for its remarkable ability to produce C2+ products, primarily due to the increased difficulty of *CO desorption, with a relatively lower energy barrier for C-C coupling (J. Am. Chem. Soc. 2022, 144, 259-269; Nat. Commun. 2024, 15, 1749). Moreover, Cu(111) favors the formation of C2+ products through the formation of *COOH and *CO (Energy Environ. Sci. 2022, 15, 3968-3977), while Cu(110) shows an intermediate product selectivity (Surf. Sci. 1995, 335, 258-263). The Cu(100) surface not only facilitates the adsorption of CO2, but also maintains an optimal binding strength with *CO, which might enhance *CO surface coverage, thereby promoting efficient coupling reactions. With regard to NO3RR (Angew. Chem. Int. Ed. 2023, 62, e202303327; ACS Catal. 2021, 11, 14417-14427), both Cu(100) and Cu(110) exhibit a notable capacity to reduce the adsorption energy of NO3-. The subsequent step (*NO3 + H+ → *NO3H) shows significantly enhanced thermodynamic favorability on Cu(100) compared to Cu(111). However, as the reaction proceeds, Cu(111) and Cu(100) remarkably facilitate the hydrogenation of *NO2 into *NOH with significantly higher efficiency compared to Cu(110), particularly on the Cu(111). Notably, *NOH was a critical N-intermediate for the first C-N coupling.\nIn summary, the integration of our DFT results with previous investigations on Cu facets revealed that both Ag(111)-Cu(111) and Ag(111)-Cu(100) exhibited favorable properties for urea synthesis. Notably, Ag(111)-Cu(100) emerged as the most promising candidate due to its reaction sites with enhanced activity, which effectively promoted the adsorption and activation of reactants, and also exhibited remarkable capability in generating and stabilizing key intermediates (*CO and *NOH), while presenting the lowest free energy barrier for two C-N couplings. These advantages would collectively enhance its performance of urea synthesis. Consequently, we predicted the urea synthesis performance ranking as follows: Ag(111)-Cu(100) > Ag(111)-Cu(111) > Ag(111)-Cu(110). However, it is important to note that the higher surface energy of Cu(100) also makes its formation more challenging under normal conditions (Nat. Commun. 2024, 15, 1749). To expose specific crystal facets, researchers have developed innovative strategies (Nat. Commun. 2022, 13, 3158; Nat. Commun. 2022, 13, 3158). Inspired by the insightful suggestion from the reviewer, in the future, we plan to investigate the design of catalysts featuring Cu(100) or Cu(100)-Cu(111) interfaces, with the aim of thoroughly exploring their potential performance.\nFigure R22. Free energy diagram of urea synthesis on Ag(111)-Cu(111), Ag(111)-Cu(100) and Ag(111)-Cu(110).\nFigure R23. The top and side views of the Ag(111)-Cu(111) model, Ag(111)-Cu(100) model and Ag(111)-Cu(110) model. The blue and red balls represent Cu and Ag atoms, respectively.\nTable R3. The free energy of adsorbent intermediates during urea synthesis on Ag(111)-Cu(111), Ag(111)-Cu(100), and Ag(111)-Cu(110).\nAg(111)-Cu(111) Ag(111)-Cu(100) Ag(111)-Cu(110)\nSystem Free Energy (eV) Free Energy (eV) Free Energy (eV)\nCO2 + NO3- 0 0 0\n*CO2 + NO3- 0.4477 0.2014 0.3912\n*CO2 + *NO3 1.3343 0.7825 0.5843\n*CO2 + *NOH -4.8814 -5.2336 -7.5700\n*CONOH -5.5087 -6.1317 -6.0259\n*CONH2 + *NO2 -13.2716 -13.1395 -15.3799\n*NOCONH2 -13.6563 -14.2358 -16.2525\n10. How was the 0.5 M KNO3 concentration determined to be optimal for flow cell experiments? Experiments can be conducted at different concentrations.\nAuthor response: We are deeply appreciative of the reviewer's insightful suggestion that experiments conducted under varying KNO3 concentrations are crucial. In response, we have supplemented the results for Ag67Cu33 at -0.52 V under different KNO3 concentrations in the flow cell, as illustrated in Figure S18. Initially, the electrolyte containing 0.03 M KNO3 (about 2000 ppm) was utilized, while it resulted in suboptimal urea synthesis performance, even inferior to that observed in an H-cell. Previous studies had demonstrated that flow cells could significantly enhance the carbon dioxide reduction reaction (CO2RR) (Adv. Mater. 2024, 36, 2303902), a finding also supported by our LSV curves (Figure 2b and 3b). Consequently, a higher concentration of KNO3 was deemed necessary to facilitate the C-N coupling reaction in the flow cell. As depicted in Figure S18b, current densities progressively increased with rising KNO3 concentrations. Furthermore, as the concentration of KNO3 increased, Ag67Cu33 exhibited progressively higher urea yield rates along with enhanced Faradaic efficiencies (FE), achieving optimal performance at 0.5 M KNO3. However, further increases in KNO3 concentrations led to a decline in urea synthesis performance, which could be attributed to excessively high KNO3 concentrations inhibiting CO2 adsorption by promoting NO3- adsorption (Angew. Chem. Int. Ed. 63, e202406515), thereby favoring competitive nitrate reduction reaction (NO3RR) and consequently reducing both urea yield rate and FE. Therefore, we selected 0.5 M KNO3 for subsequent electrochemical tests in the flow cell. Detailed analysis and discussion of these data are also provided in the main text, specifically from line 243 on page 7 of the revised manuscript.\nFigure S18.(a) The actual test process of a three-electrode flow cell. The i-t curves(b), UV-vis absorption spectra for urea(c), together with the corresponding urea yield rate and Faradaic efficiency(d) of Ag67Cu33 at -0.52 V under varying KNO3 concentrations."}, {"round": 2, "reviewer_comments": {"reviewer_1": "The authors have addressed all my concerns and thus I would suggest the acceptance of this manuscript as is.", "reviewer_2": "The authors have well addressed my issues and this work can be published now.", "reviewer_3": "The manuscript was significantly improved after the revision. Authors have answered all the questions and made necessary amendments. The present version of the manuscript can be accepted after some minor revision.\n(1) The authors claim, \"Ag67Cu33 exhibited excellent bi-functional electrosynthesis for urea (FE: 56.07%, yield: 104.6 mmol h-1 gcat.-1) and HCOOH (FE: over 90%) at 40 mA cm-2 in a two-electrode system\". They should provide the NMR peaks of HCOOH after the electrolysis for better understanding.\n(2) In Figure 5b, the formaldehyde oxidation reaction with the synthesis of HCOOH. If HCHO in the presence of OH– produces HCOOH, what is the role of CO2 and nitrate during the synthesis of HCOOH? They should check the LSV profile for HCHO but in the absence of CO2 and nitrate.\n(3) The authors should provide the yield rate and FE equations for formic acid calculations."}, "author_response": "We extend our heartfelt gratitude for your meticulous review and insightful suggestions regarding our manuscript. Your professional feedback offered invaluable guidance, which substantially enhanced both the quality and depth of this research. Please accept our sincere appreciation for your expert contributions to the advancement of this research.\nWe are deeply appreciative of the reviewer's professional and insightful suggestions, particularly those related to the theoretical computational methodology. Your meticulous feedback not only has substantially improved the quality of our work but also holds significant guiding value for our future research in this field.\nWe sincerely appreciate the reviewer's insightful and professional feedback, which has substantially enhanced the quality of our manuscript. We also express our gratitude for recognizing our improvements. Your latest suggestions have highlighted the omission of critical information regarding the two-electrode system, which we deeply acknowledge. In accordance with your astute recommendations, we have integrated essential details into our revised manuscript.\nBelow, we will offer thorough and detailed responses to each of your valuable comments:\n(1) The authors claim, \"Ag67Cu33 exhibited excellent bi-functional electrosynthesis for urea (FE: 56.07%, yield: 104.6 mmol h-1 gcat.-1) and HCOOH (FE: over 90%) at 40 mA cm-2 in a two-electrode system\". They should provide the NMR peaks of HCOOH after the electrolysis for better understanding.\nAuthor response: We sincerely express our gratitude for your timely and insightful reminder. Guided by your astute suggestion, we have carefully incorporated the relevant NMR peaks of HCOOH into the revised supplementary information, significantly improving its comprehensibility and clarity. Please refer to the section titled \"Determination of HCOOH and CH3OH\" in the revised manuscript (page 29) for further details.\nFigure S53. The E-t curves(a), the corresponding UV data for urea(b) and NMR data for HCOOH(c) obtained on Ag67Cu33 in the two-electrode system ranging from 15 mA cm-2 to 45 mA cm-2.\nFigure S55. The urea yield rate and corresponding FE in the cathode(a) and HCOOH yield rate and corresponding FE in the anode(b) of Ag67Cu33 in the two-electrode system at 40 mA cm-2 by 10 cycles, with each cycle presenting 1 h. The corresponding UV data for urea(c) and NMR data for HCOOH(d).\nFigure S57. The E-t curves(a), the corresponding UV data for urea(b) and NMR data for HCOOH(c) obtained on Ag67Cu33 in the two-electrode system at 40 mA cm-2 of long-term electrolysis for 60 h.\n(2) In Figure 5b, the formaldehyde oxidation reaction with the synthesis of HCOOH. If HCHO in the presence of OH– produces HCOOH, what is the role of CO2 and nitrate during the synthesis of HCOOH? They should check the LSV profile for HCHO but in the absence of CO2 and nitrate.\nAuthor response: We extend our deepest gratitude to the reviewer for their insightful and expert recommendations. In Figure 5, we utilized a two-electrode flow cell for further investigation. The Formaldehyde Oxidation Reaction (FOR) occurred at the anode (HCHO + 2OH- → HCOO- + 1/2H2 + H2O + e-), while the urea synthesis process was carried out at the cathode (NO3- + CO2 + 18 H+ + 16e- → NH2CONH2 + 7 H2O). Both reactions were efficiently catalyzed by the Ag67Cu33 catalyst, and the two reaction systems were separated using an anion exchange membrane (Figure R1). In accordance with the reviewer's suggestion, we conducted Linear Sweep Voltammetry (LSV) tests in a system where the anolyte consisted of 1 M KOH with 0.6 M HCHO, and the catholyte was composed of 0.5 M K2SO4, free of CO2 and NO3-. As illustrated in Figure 5b, this system (FOR||HER) demonstrated a lower current density compared to the FOR||Urea synthesis system (anolyte: KOH solution with HCHO, catholyte: CO2 and NO3-), primarily attributed to its limited HER capability. Thus, the success of our objective to reduce the potential of the two-electrode system and produce high-value chemicals was validated by coupling appropriate reactions at both electrodes (Energy Environ. Sci., 2025,18, 2804-2816).\nFurthermore, the spontaneous disproportionation reaction of HCOH (Cannizzaro Reaction) in the presence of OH-, leading to the formation of HCOOH, was taken into account in the calculation of the HCOOH yield rate and Faradaic Efficiency (FE) produced via Ag67Cu33-catalyzed FOR at the anode. While the explanation about it was inadvertently overlooked in our initial manuscript. In the revised manuscript, we have incorporated additional details to comprehensively describe this (please refer to line 25 on page 29). Based on the chemical equation (Cannizzaro Reaction: 2HCHO + OH− → CH3OH + HCOO−), it is evident that one molecule of HCOOH forms spontaneously alongside one molecule of CH3OH. By quantifying the amount of CH3OH generated, we can deduce the quantity of HCOOH formed via the spontaneous decomposition of HCOH. This information allows us to accurately calculate the amount of HCOOH produced through the FOR process catalyzed by Ag67Cu33. It is important to note that commercial HCOH solutions inherently contain a predetermined concentration of CH3OH, which serves to inhibit the spontaneous polymerization of HCOH. Therefore, pre-existing CH3OH content must be measured beforehand. To determine the precise HCOOH content derived from electrocatalytic FOR, the following formula should be applied:\nwhere C_HCOOH-t (mmol L-1) represents the concentration of HCOOH and CH3OH in the post-electrolyte as determined by 1H NMR analysis, and C_CH3OH-0 denotes the initial concentration of CH3OH in a 0.6 M HCOH solution (without any treatment).\nFigure R1. Scheme of the two-electrode flow cell for coupling cathodic urea synthesis on Ag67Cu33 with anodic FOR on Ag67Cu33, separated by an anion exchange membrane.\nFigure 5b. LSV curves for various conditions in the two-electrode flow cell.\n(3) The authors should provide the yield rate and FE equations for formic acid calculations.\nAuthor response: We extend our sincere gratitude to the reviewer for drawing our attention to the omission of critical calculation formulas, which are indispensable for enhancing readers' understanding. To determine the yield rate and Faradaic Efficiency (FE) of HCOOH in a two-electrode flow cell, we utilized the 1H NMR technique (Figure S29) to accurately measure the concentrations of HCOOH and CH3OH in the post-electrolyte (a detailed explanation of the calculation process was provided in Question 2). Thereafter, the concentration of HCOOH generated via Ag67Cu33-catalyzed FOR at the anode was calculated using the following formula:\nwhere C_HCOOH-t (mmol L-1) represented the concentration of HCOOH and CH3OH in the post-electrolyte as determined by 1H NMR analysis, and C_CH3OH-0 denoted the initial concentration of CH3OH in a 0.6 M HCOH solution (without any treatment).\nSubsequently, the yield rate and Faradaic efficiency of HCOOH achieved using the Ag67Cu33 catalyst via the FOR process were calculated based on the following equations (Angew. Chem. Int. Ed. 2023, 62, e202302950):\nwhere C_HCOOH (mmol L-1) represented the concentration of HCOOH produced by Ag67Cu33 catalyst at the anode, V denoted the volume of the electrolyte solution, t indicated the reaction time, m referred to the mass of the catalyst, F was the Faraday constant (96485.34), and Q signified the applied electrical charge. For a more detailed comprehension, please consult page 29 of the revised manuscript.\nWe extend our sincerest gratitude for the significant time and effort you have dedicated to reviewing our work. Your meticulous and insightful feedback has offered us invaluable guidance. We sincerely hope that our comprehensive response not only addresses your concerns but also meets your expectations. Lastly, we once again express our deepest appreciation for your contributions."}, {"round": 3, "reviewer_comments": {"reviewer_3": "The manuscript has improved significantly after revision. The authors have answered all questions and made the necessary corrections. The current version of the manuscript is acceptable for this journal."}, "author_response": ""}]

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