Membrane-free CO_2 hydrogenation electrolyzer for salt precipitation management in acidic electrochemical CO_2 reduction

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

Abstract
Electrochemical CO2 reduction (ECR) in acidic electrolytes minimizes CO2 loss and carbonate formation issues, allowing for high CO2 utilization efficiency and showing good potential for practical CO2 upgrading applications. However, in the membrane-based electrolyzer, the proton transfer efficiency across the membrane from the anolyte to the catholyte is crucial for the stability of the catholyte pH in acidic ECR, especially at high current density and during long-term electrolysis. Here, we investigate the effects of proton transfer efficiency and salt precipitation in different acidic ECR electrolyzer and propose a membrane-free CO2 hydrogenation electrolyzer, which couple CO2 reduction and hydrogen oxidation. This electrolyzer design effectively maintains a stable electrolyte pH during long-term electrolysis, and simultaneously achieves high Faradaic efficiency for HCOOH production, high single-pass carbon utilization efficiency, and a lower cell voltage. At a current density of 100 mA cm−2, the system requires only 1.7 V to achieve a 90% HCOOH Faradaic efficiency and demonstrates stable operation for 208 hours.

Introduction
Electrochemical CO2 reduction (ECR) powered by low-cost renewable electricity offers a promising route for producing valuable chemicals and feedstocks1–3, such as formic acid4–7, carbon monoxide8, and ethylene9,10, while simultaneously mitigating greenhouse gas emissions9–13. Although high selectivity and activity for ECR have been achieved in alkaline9,14 and neutral electrolytes15,16, the formation and precipitation of carbonates impede the implementation of efficient and stable CO2 conversion systems17–21. Specifically, during ECR, one CO2 molecule reacts with two hydroxide ions (OH−) to form one equivalent of carbonate (CO32−) for every two electrons transferred22. Consequently, the carbon utilization efficiency, defined as the fraction of the CO2 feed converted to ECR products, is critically low (less than 50%) due to the majority of CO2 reacting with hydroxide ions to form carbonates23. The recovery of CO2 from a carbonate electrolyte is an energy-intensive process22. Furthermore, after long-term electrolysis, the precipitation of carbonates, such as potassium carbonate, can block CO2 diffusion, thereby inhibiting ECR23–25. In addition, for certain products such as formic acid (HCOOH) and acetic acid, electrolysis in alkaline or neutral electrolytes presents an additional challenge: the acids cannot be directly obtained because their corresponding anions exist in salt form26,27, leading to high separation costs28.

ECR in acidic electrolytes offers an avenue to address the carbonate formation issues8,29. Specifically, when H3O+ is the proton source for ECR, no OH− is generated17,30. Additionally, when H2O is the proton source, any locally generated carbonate reacts with H+ in the bulk electrolyte, converting it back to CO2 31,32. Moreover, acid products can be readily recovered by distillation in acidic conditions23. Efforts have been dedicated to developing electrolytes and catalytic surfaces that minimize the competing hydrogen evolution reaction (HER) and increase ECR selectivity, activity, and stability32,33. For instance, introducing alkali cations (such as K+) and functionalizing electrocatalysts with cationic molecules can alter the local environment around the electrocatalysts, promoting acidic ECR performance17,34–38. However, the membrane39, an essential component in ECR, which crucially impacts the composition of the catholyte and thus affects acidic ECR performance, has received less attention. In acidic ECR electrolyzers, a cation exchange membrane (CEM), such as Nafion 117 and Nafion 211 membrane, is commonly used, allowing proton transport from the anode to the cathode and avoid the oxidation of liquid ECR products in the anode side17,34–37. The use of a CEM introduces another often-overlooked issue: the proton transport efficiency from the anolyte to the catholyte can influence the pH of both compartments, potentially leading to salt precipitation over long-term electrolysis.

Here, we explored the crucial role of the membrane in acidic ECR. In systems using a cation exchange membrane, the catholyte pH was found to increase progressively during electrolysis23. To address the limited H+ transport caused by the membrane, a membrane-free electrolyzer was developed, effectively minimizing the pH increase of the electrolyte. To prevent oxidation of ECR products at the anode, we utilized hydrogen oxidation reaction to replace the conventional oxygen evolution reaction, which operates at a much lower oxidation potential than that of formic acid. Coupled with a Bi0.5-Ag0.5 catalyst engineered for fast charge transfer, the membrane-free CO2 hydrogenation electrolyzer enabled stable HCOOH production for 208 h, maintaining over 90% Faradaic efficiency (FE) at 1.7 V in acidic electrolyte at 100 mA cm−2.

Results
Membrane H+ transfer efficiency impact catholyte pH
In acidic ECR, the catholyte pH typically ranges from 0.5 to 4, exhibiting negligible CO2 loss17. This is evident from the pH-dependent distribution of CO2, bicarbonate (HCO3−), and CO32− (Supplementary Fig. 1)40. The fraction of HCO3− ranges from 10−6 to 10−2 when the solution pH is between 0.42 and 4.38. However, as the pH continues to increase, the fraction of bicarbonate and carbonate increases exponentially, eventually becoming the dominant species and resulting a carbonate formation issues at the cathode. Therefore, it is crucial to maintain stable catholyte pH, preferably without any fluctuations, to effectively address carbonate formation issues, especially in long-term and/or high current practical applications.

The acidic CO2 electrolysis system typically employs a flow-type reactor (Fig. 1a), where a certain amount of anolyte and catholyte are recycled at a fixed flow rate. The anolyte and catholyte consist of different concentrations of protons (H+), alkaline metal cations (M+), and anions (A−). During CO2 electrolysis, H+ and/or M+ can crossover from the anolyte to the catholyte driven by the electrical field and diffusion and then influence the catholyte and anolyte H+ and M+ concentration. To better describe the influence of H+ transport, we introduce a parameter, ηH+, to better describe the H+ transfer effects in acidic ECR, which is highly correlated with the overall H⁺ flux. The ηH+, defined as the ratio of the transferred H+ to the transferred electron, is an electrolysis charge-normalized parameter used to quantify the H⁺ transfer associated with electron transfer process. The ηH+ directly influences the catholyte pH during acidic ECR and is affected by electrolysis time, applied current, and the membrane properties. The relationship between the catholyte pH, ηH+, electrolysis current and time can be easily obtained according to the following Eq. (1):

ηH+ = Transfered H+ / Passedelectron × 100% = Transfered H+ / (Current × Time) = F × NA × 100% (1)

where F is the Faraday constant, NA is Avogadro constant. Assuming an initial catholyte pH of 2 and a volume of 10 mL, Fig. 1b shows that a constant electrolysis current of 200 mA, the catholyte pH is influenced by both ηH+ and electrolysis time. Even a smaller decrease of ηH+ lead to a significant increase in catholyte pH. Similarly, as shown in Fig. 1c, when the electrolysis time is fixed at 10 min, the catholyte pH increase sharply with the decrease of ηH+ or increase of the electrolysis current.

From aforementioned analysis, we found that the ηH+ significantly influences the catholyte pH, leading to issues with carbonate formation and related challenges associated with fluctuations in the catholyte pH (Fig. 1d).

When ηH+ is 100%, the pH of both the catholyte and anolyte will remain at their initial states. However, achieving a 100% ηH+ is challenging according to the physicochemical principles that govern ion transport. The overall flux of arbitrary ion i through a membrane, Ji, consists of three terms: the diffusion flux (Ji,diffusion), caused by the chemical potential gradient; the electric transference (Ji,electric), caused by the electric potential gradient; and the convection flux Ji,convection, associated with the bulk velocity41,42. The overall flux can be expressed by the following Eq. (2):

Ji = Ji,diffusion + Ji,electric + Ji,convection = −D ∇Ci + ziCi F/RT ∇φ + Ci v (2)

where D, Ci, zi are the diffusion coefficient, concentration, and valence of the arbitrary ion i, respectively. F is the Faraday constant, R is the gas constant, T is temperature, φ is the electrical potential, and v is bulk fluid velocity. For homogenous and nonporous membranes, the convection term can be neglected. This relation is known as the Nernst-Planck equation and applies to all mobile ions present34. In an acidic ECR system, several factors influence the overall Ji. Firstly, the properties of the membrane, such as its chemical structure, water content, pore size, will affect the D and thus affect the Ji,diffusion. Membranes with higher Di facilitate H+ transport and help minimize the pH gradient between the anode and cathode. Secondly, the composition of the anolyte and catholyte, including the anolyte species (with or without alkaline cations), concentration, as well as the concentration gradient between the anolyte and catholyte, also influence Ji,diffusion. For example, the transport of competing cations can hinder H+ transport, while a high H+ concentration at the anode enhances H+ transport due to a steeper concentration gradient. Thirdly, the applied potentials affect the Ji,electric. A higher voltage between the anode and cathode can enhance H+ transport from the anode to the cathode. Consequently, the overall Ji will directly influence the pH changes of the catholyte and anolyte, and finally influence the acidic ECR performance.

We then employed COMSOL to further analyze the influence of membrane and electrolyte composition to proton transport (Supplementary Fig. 2 and Methods). We firstly investigated the influence of membrane thickness (Supplementary Fig. 3), a thicker membrane restricted proton transfer, resulting in a non-uniform proton flux and hindering transport from the anode to the cathode (right to left). Consequently, the proton concentration distribution across the cell was non-uniform, and the ηH+ at the cathode side was low. In contrast, in the absence of a membrane (bottom panel), proton flux was faster and more uniform from the anode to the cathode, leading to a uniform proton concentration distribution and a significantly higher ηH+ near the cathode compared with the system with membrane. In addition, higher proton selectivity could also promote proton transport from the anode to the cathode (Supplementary Fig. 4). As proton selectivity decreased, H+ flux, H+ concentration, and ηH+ at the cathode side all declined, indicating that a membrane with high proton selectivity is more effective at maintaining catholyte pH. We also investigate the influence of Electrolyte composition (Supplementary Fig. 5). As H+ concentration increases, proton transport from the anode to the cathode becomes more difficult. When both the anode and cathode have high proton concentrations, reduced chemical potential, strong ionic shielding, and near-saturation of proton-conducting channels in the membrane collectively limit the effective proton migration rate. Taken together, a thinner membrane with high proton selectivity and an electrolyte with low proton concentration are more favorable for proton transport.

ECR catalyst design and evaluation of the impact of H+ transport efficiency on ECR performance
To support the above analysis, we evaluated the ECR performance in a conventional ECR and OER paired electrolyzer with cation exchange membrane (ECR//CEM//OER), as shown in Fig. 2a. Nafion 117 and Pt/C were employed as the CEM and OER catalysts, respectively. In the initial electrolyzer configuration, the distance between the anode and cathode was set to 10 mm. CO2-to-HCOOH conversion was selected as the model reaction due to the high economic value of HCOOH. We selected Bi-based materials as the catalyst due to it’s low toxicity and high abundance, which also exhibit high formate selectivity43. However, the electrical conductivity of Bi is relatively low, leading to the higher cell voltage. Considering this, we introduced high electrical conductive silver into the Bi-based materials to enhance the charge transfer and lower the ECR potentials44. We synthesized different composed Bi-Ag and pure Bi through grinding method. According to the different precursor stoichiometric ratios, the catalysts are designated as Bi, Bi0.3-Ag0.7, Bi0.5-Ag0.5, Bi0.7-Ag0.3, respectively. Scanning electron microscopy (SEM) and transmission electron microscopy (TEM) images show that Ag nanoparticles are dispersed uniformly on the Bi nanosheets in the as-prepared samples and EDS mapping demonstrate the uniform distribution of Bi and Ag within the structure (Fig. 2b, c, Supplementary Figs. 6-10). X-ray diffraction (XRD) patterns confirm the co-existence of metallic Bi and Ag in the structure for all Bi-Ag samples (Fig. 2d). X-ray photoelectron spectroscopy (XPS) also confirmed the existence of Bi and Ag (Supplementary Fig. 11). Bi0.5-Ag0.5 electrocatalyst with efficient charge transfer capability, exhibiting high activity, selectivity, and stability for CO2-to-HCOOH conversion was designed (details for electrocatalysts design are discussed in Supplementary Note 1, Supplementary Figs. 12-14). The FEHCOOH remained above 90% across a current density range of 100−400 mA cm−2, with a HCOOH partial current density reaching 364 mA cm−2 at −1.31 V vs. RHE (reversible hydrogen electrode) in acidic electrolyte (Fig. 2e, f).

After ECR, we disassembled the electrolyzer and characterized the catalysts post-reaction. The morphology of Bi0.5-Ag0.5 was well preserved (Supplementary Fig. 15). XPS analysis revealed electron transfer between Bi and Ag: the Bi 4f core levels exhibited a positive binding-energy shift upon introducing Ag (Supplementary Fig. 16), indicating electron withdrawal from Bi. This electronic structure change is further supported by the product distribution: although Ag typically favors CO formation over formate, only trace CO was detected in our system, suggesting that the electronic structures of both Bi and Ag were altered. To probe the reaction intermediates, in-situ attenuated total reflection surface-enhanced infrared absorption spectroscopy (ATR-SEIRAS) was performed (Supplementary Fig. 17). A characteristic band at approximately 1400 cm−1 was observed on both Bi and Bi0.5-Ag0.5, commonly assigned to *OCHO intermediates adsorbed on the catalyst surface45. This is consistent with previous reports that formate is generated through hydrogenation of the C atom in *OCHO46. As the potential shifted negatively, the *OCHO signal intensity gradually increased, indicating higher surface coverage of this intermediate. Notably, Bi0.5-Ag0.5 exhibited a pronounced increase in the intensity of a band near 1250 cm−1 with increasing overpotential, attributed to the asymmetric −COO− stretching vibration of activated *CO2− species47. In contrast, pure Bi showed a much weaker increase, suggesting that Bi0.5-Ag0.5 is more favorable for CO2 adsorption and activation. In-situ shell isolated nanoparticle-enhanced Raman spectroscopy (SHINERS) measurements (Supplementary Fig. 18) further confirmed that both Bi and Bi0.5-Ag0.5 follow similar *CO2− pathways toward formate formation48,49, but Bi0.5-Ag0.5 required a lower overpotential for *CO2− formation. These results indicate that the improved CO2-to-formate conversion performance of Bi0.5-Ag0.5 arises from both electronic effects that promote CO2 adsorption/activation and the enhanced charge-transfer process enabled by the high electrical conductivity of Ag. The ECR performance of the Bi0.5-Ag0.5 catalyst was also assessed in neutral and alkaline electrolytes, where formate Faradaic efficiencies consistently exceeded 90% (Supplementary Fig. 19). Note that the electrolysis time was controlled to 10 min to exclude the possible effects of salt precipitation.

We then assessed the long-term ECR performance of the ECR//CEM//OER electrolyzer at 100 mA cm−2 (Fig. 2g). The pH of the catholyte was monitored in real time using a pH meter throughout electrolysis. During the first 5 hours, the pH rapidly increased to 5 due to the fast consumption of protons, followed by a slower rise over the next 3 hours, eventually stabilizing at around 7.2. The initial ηH+ was only ~84% during the first hour, leading to the sharp increase in catholyte pH. The gradual improvement in ηH+ may be attributed to the growing pH gradient between the anolyte and catholyte, which likely enhanced Ji,diffusion. The average FEHCOOH declined from 81% to 75% after 8 hours of operation.

To investigate the mechanism of performance degradation, we characterized the electrode after ECR. As shown in Fig. 2h, XRD analysis confirmed the presence of KHCO3 and K2CO3 precipitates within the gas diffusion layer. Energy-dispersive X-ray analysis (EDS) elemental mapping was then conducted to visualize the spatial distribution of (bi)carbonate species across the gas diffusion electrode (GDE) cross-sections. The K+ penetration depth reached as deep as ~76 µm (Fig. 2i). Inductively coupled plasma optical emission spectroscopy (ICP-OES) analysis revealed a K content of 4.81% in the electrode after 8 hours electrolysis in ECR//CEM//OER electrolyzer (Supplementary Table 3). Additionally, a Raman peak at 1055 cm−1 further confirmed the formation of K2CO3 (Supplementary Fig. 20)50. The results validated the assumption that the inefficient ηH+ in the traditional ECR//CEM//OER electrolyzer led to a rise in catholyte pH from acidic to neutral, which promoted (bi)carbonate precipitation on the cathode surface, subsequently blocking CO2 diffusion and suppressing ECR activity.

Electrolyzer design to enable stable catholyte pH
The ηH+ and stability of catholyte pH are crucial for acidic ECR. A membrane-free electrolyzer provides an alternative that bypasses the limitations in H+ transfer from the anolyte to the catholyte caused by the membrane. We then designed the ECR and OER paired electrolyzer without membrane (ECR//OER, Fig. 3a). As expected, the electrolyte pH maintained nearly unchanged during the long-term test and the ηH+ was close to 100% (Fig. 3b), owing to the rapid neutralization of hydroxide ions formed at the cathode and hydrogen ions generated at the anode through electrolyte flow. These results highlight the advantages of a membrane-free electrolyzer compared to a membrane-based electrolyzer. However, the average FEHCOOH decreased to 54% after 8 h electrolysis (Fig. 3b), possibly due to the oxidation of HCOOH at the anode, as the equilibrium potential for formic acid oxidation reaction (FAOR, 0.2 V vs. RHE)51 is lower than the standard potential for OER (1.23 V vs. RHE)4,52.

Motivated by recent HOR-coupled electrosynthesis systems for nitrogen reduction53,54, we developed a membrane-free CO2RR−HOR MEC configuration that similarly utilizes the thermodynamic advantages of hydrogen oxidation to reduce the cell voltage. Compared to OER (1.23 V vs. RHE), hydrogen oxidation reaction has a much lower standard equilibrium potential (0 V vs. RHE, Supplementary Table 4, Supplementary Fig. 21), which is also lower than the standard equilibrium potential for HCOOH oxidation (0.2 V vs. RHE). Inspired by pioneering work21 on bicarbonate reduction for CO production using membrane-free electrolyzers coupled with HOR in the anode, we coupled HOR with ECR in a membrane-free CO2 hydrogenation electrolyzer (ECR//HOR, Fig. 3c) to address the issue of HCOOH oxidation at the anode54 caused by the absence of a membrane (Supplementary Fig. 22). As shown in Fig. 3d, the cell voltage of the ECR//HOR electrolyzer (3 V) is lower than that of the ECR//OER electrolyzer at a current density of 100 mA cm−2 with the same electrolyte composition and anode to cathode distance. The pH of the electrolyte remained nearly constant, around 1.8, during 8 hours of electrolysis, and the FEHCOOH in the ECR//HOR electrolyzer stayed above 90%, highlighting the advantages of the membrane-free CO2 hydrogenation electrolyzer.

For comparison, the ECR performance and the pH change trend in an ECR and HOR paired electrolyzer with a membrane was tested (ECR//CEM//HOR, Fig. 3e). The cell voltage increased to ~ 5.2 V at a current density of 100 mA cm−2 due to the membrane resistance55 (Fig. 3f). The catholyte pH increased to approximately 7.8 after 8 h of electrolysis, which is even higher than that of ECR//CEM//OER. This is possibly due to the higher ECR//CEM//OER cell voltage, which creates a stronger electric field between the anode and cathode, thereby enhancing proton diffusion across the membrane and leading to a lower catholyte pH. In addition, similar to the ECR//CEM//OER electrolyzer, the FEHCOOH continued to decrease throughout the 8-hour test, likely due to the same salt precipitation issues.

We then conducted a detailed post-analysis of the cathode electrode after 8 h of electrolysis for the three types of electrolyzers mentioned above. As shown in Fig. 3g, the K+ penetration depth from EDS and K+ content from ICP-OES results are highly related to the catholyte pH after electrolysis. For the ECR//HOR electrolyzer, only trace amounts of KHCO3/K2CO3 salt were detected (Fig. 3h), and the penetration depth of K was approximately 18 µm (Fig. 3i). This was consistent with the good performance of the membrane-free ECR//HOR electrolyzer, as less salt precipitation formed during long-term electrolysis. Compared to the ECR//HOR electrolyzer, the ECR//OER electrolyzer exhibited greater K+ penetration and content (Supplementary Fig. 23), likely due to stronger electric field effects induced by the higher cell voltage. Similarly, the deeper K+ penetration observed in the ECR//CEM//OER electrolyzer compared to the ECR//CEM//HOR configuration can be attributed to the same factor. As for the K content, the lower K content in ECR//CEM//OER compared to ECR//CEM//HOR is possibly due to the lower pH of the ECR//CEM//OER catholyte after electrolysis (Supplementary Fig. 24). Lower pH favors the formation of KHCO3, while higher pH promotes the formation of K2CO3, which contains a higher K content than KHCO3. Raman spectra of the cathode electrode after ECR further confirmed our analysis that (bi)carbonate species formed in the membrane-based electrolyzer, while the membrane-free electrolyzer exhibited significantly lower signal intensity (Supplementary Fig. 25).

Electrolyzer optimization towards practical applications
To evaluate the practical applicability of our membrane-free CO2 hydrogenation electrolyzer, we measured the single-pass conversion efficiency (SPCE) of the ECR//HOR electrolyzer at a current density of 400 mA cm−2 under varying CO2 flow rates (Supplementary Fig. 22c). A SPCE of 77% was achieved at a CO2 flow rate of 3 sccm, which surpasses that of other electrolyzers (Fig. 4a). We then conducted a technoeconomic analysis (TEA) based on the results of the 8-h long-term electrolysis test. The TEA revealed that the production cost of HCOOH is significantly reduced in the membrane-free ECR//HOR CO2 hydrogenation electrolyzer (Fig. 4b, Supplementary Note 2), owing to the excellent system performance (Fig. 4c) and the elimination of membrane-related costs.

To further reduce product cost, we designed an electrolyzer with a reduced electrolyte flow layer thickness, decreasing the distance between the anode and cathode from 10 mm to 2 mm, which effectively lowers cell resistance and thus the cell voltage (Supplementary Fig. 26). As a result, the cell voltage of the ECR//HOR electrolyzer decreased from 3 V to approximately 1.7 V at 100 mA cm−2 (Fig. 4d). The optimized system exhibited stable operation for 50 hours at 100 mA cm−2, maintaining a FEHCOOH of 90%. It continuously produced HCOOH at a rate of 0.008 mol L−1 h−1 (Supplementary Fig. 27), and the maximum HCOOH concentration reached 0.1 M, demonstrating the practical potential of our design. Post-analysis of the cathode electrode after 50 hours revealed minor salt precipitation (Fig. 4e, Supplementary Fig. 28), indicating that the optimized membrane-free ECR//HOR electrolyzer effectively mitigates severe carbonate accumulation even under prolonged operation. However, the minor salt precipitation obstructed the gas flow tubing. To address this issue, we used acidic humidification strategy56, which enabled stable CO2 reduction for over 208 hours (Fig. 4f, Supplementary Table 5). Furthermore, long-term electrolysis conducted in a 4 cm2 flow electrolyzer (2 cm × 2 cm electrode area) also demonstrated stable performance, maintaining a HCOOH FE of approximately 90% over a 48-h operation period (Supplementary Figs. 29-30). To highlight the advantage of the membrane-free acidic ECR//HOR configuration, we evaluated its performance against membrane-based systems using neutral (0.5 M K2SO4) and alkaline (1 M KOH) electrolytes (Supplementary Fig. 31). Although all systems showed similar formate selectivity (~90%), the membrane-free acidic setup exhibited a much lower cell voltage (1.7 V) compared to the neutral (2.2 V) and alkaline (2.7 V) systems, demonstrating its high energy efficiency for CO2 electroreduction. Morevoer, we performed electrochemical measurements in ECR//CEM//OER electrolyzer using the same 2-mm flow channel and observed high cell voltage and a notable decrease in HCOOH FE compared to the ECR//HOR system, primarily due to severe carbonate accumulation at the cathode (Supplementary Fig. 32). In addition, to evaluate potential CO crossover and poisoning at the anode during long-term electrolysis, CO stripping was performed on the Pt/C anode after stability tests. No CO oxidation peak was observed on the electrode after stability test, indicating that no significant CO adsorption occurred on the Pt/C anode during long-term operation (Supplementary Fig. 33). Together with the stable low cell voltage (~1.7 V at 100 mA cm−2, Fig. 4d, f), these results indicate that the CO generated at the cathode has a negligible impact on the overall cell voltage and electrochemical performance. The H2 permeation test also indicated that H2 on the anode side scarcely crossed the electrolyte to reach the cathode (Supplementary Fig. 34). Moreover, the lower cell voltage could further reduce the HCOOH production cost to ~ 978.6 USD ton−1 (Fig. 4g).

Integrated system for efficient electron utilization
Although the ECR and HOR coupled membrane-free system can significantly reduce the cell voltage, the production cost of HCOOH remains high due to the expensive hydrogen source. While H2 production from water splitting could lower the H2 cost, the OER at the anode still requires a high potential and generates low-value oxygen as a byproduct. In contrast to OER (1.23 V vs. RHE), MOR (methanol oxidation reaction, 0.04 V vs. RHE) has a much lower thermodynamic equilibrium potential57. Additionally, by optimizing the electrocatalyst, the MOR product can be tuned to high-value formate instead of CO2 58. Therefore, we propose to couple the ECR//HOR electrolyzer with the HER//AEM//MOR electrolyzer, using the H2 produced by HER as the hydrogen source for HOR, and producing both formate and HCOOH in the two electrolyzers (Fig. 5a). Thus, electron utilization and the HCOOH production rate could be doubled.

The MOR performance was evaluated in an anion exchange membrane-based membrane electrode assembly (MEA) electrolyzer. NiFe LDH and commercial 40 wt% Pt/C were used as the anode and cathode catalysts, respectively (Supplementary Note 3, Supplementary Figs. 35-40). The FEformate maintained ~90% during a 5-hour test at 100 mA cm−2 using 1 M KOH + 1 M MeOH as the anolyte, demonstrating the good stability of this system (Supplementary Fig. 41). We then integrated the HER//AEM//MOR electrolyzer with the ECR//HOR electrolyzer (Fig. 5a, Supplementary Fig. 42). The hydrogen generated at the cathode of the HER//AEM//MOR electrolyzer was directly fed to the anode of the ECR//HOR electrolyzer as the hydrogen source, eliminating the need for an external hydrogen supply. Both the ECR//HOR and HER//AEM//MOR electrolyzers were operated at a current density of 100 mA cm−2 with a reactor area of 1 cm2. The measured cell voltages for the HER//AEM//MOR and ECR//HOR electrolyzers were 1.7 V and 1.7 V, respectively (Fig. 5b). The coupled system operated stably for 5 h, achieving a combined FE for formate and HCOOH exceeding 180% (Fig. 5c). As expected, the pH of the electrolyte in the ECR//HOR electrolyzer remained stable at ~1.8 (Supplementary Fig. 43).

A comprehensive TEA was carried out (Supplementary Note 2, Supplementary Figs. 44-52). Assuming a CO2 feedstock cost of 40 USD ton−1 and a methanol feedstock cost of 470 USD ton−1, and an electricity price of 0.03 USD kWh−1, integration with a separation process involving electrodialysis followed by extraction–distillation yields an estimated levelized production cost of formic acid of approximately 2041.1 USD ton−1 (Fig. 5d). The major cost components contributing to this value include: separation cost, electrolyzer capital cost, and electricity expense. The separation cost (1267.7 USD ton−1) accounted for approximately half of the total cost. This resulted from the low initial product concentration (0.1 M) and the presence of a KOH and formate mixture from the HER//AEM//MOR electrolyzer, which required an energy-intensive electrodialysis process to concentrate the produced formate to 2 M. In addition, because electrolysis was conducted at a low current density (100 mA cm−2), achieving the targeted production rate requires a larger electrolyzer area (electrolyzer cost59: 310 USD kW−1 at 1.7 V and 1.7 A cm−2, corresponding to 8959 USD m−2), leading to a higher capital investment. Although the production cost higher than the current market price60 (830 USD ton−1), it is expected to decrease further with reductions in electrolyzer capital cost and electricity price (Supplementary Fig. 52). In addition, compared with direct H2 supply (Supplementary Note 4), this approach reduces the costs associated with H2 transport and storage (factors not considered in our TEA) and also enhances operational safety. Future work should therefore focus on increasing the electrolysis current density to lower electrolyzer capital costs and on developing new separation methods to reduce separation costs.

Methods
Chemicals and materials
Bismuth(III) nitrate pentahydrate (Bi(NO3)3·5H2O, 98%), silver nitrate (AgNO3, 99%), sodium borohydride (NaBH4, 99%), nickel (II) nitrate hexahydrate (Ni(NO3)2·6H2O, 99%), iron (III) nitrate nonahydrate (Fe(NO3)3·9H2O, 98%), urea (98%), potassium sulfate (K2SO4, 99%), sulfuric acid (H2SO4, 95%-98%), potassium hydroxide (KOH, 90%) were purchased from Sigma-Aldrich. Carbon paper with microporous layer (YLS-30T) were purchased from Suzhou Sinero Technology Co., Ltd. All chemicals were used as received.

Synthesis of Bi xAg 1-x
For the synthesis of Bi0.5-Ag0.5, bismuth nitrate (0.076 mmol) and silver nitrate (0.076 mmol) were mixed in a certain stoichiometric ratio, ground evenly with a mortar, and then NaBH4 (0.477 mmol) was added and ground. Then, a few drops of methanol were added dropwise to remove excess NaBH4. The product after the reaction was added to a centrifuge tube, and an appropriate amount of methanol (15 mL) was added and centrifuged 3 times, and then dried in a vacuum drying oven. Similarly, for the preparation of Bi, Bi0.3-Ag0.7, and Bi0.7-Ag0.3 samples, the total amount of metal precursors was fixed at 0.152 mmol. For the Bi sample, 0.152 mmol of Bi(NO3)3·5H2O was used. For the Bi0.3-Ag0.7 sample, 0.045 mmol of Bi(NO3)3·5H2O and 0.107 mmol of AgNO3 were employed. For the Bi0.7-Ag0.3 sample, 0.107 mmol of Bi(NO3)3·5H2O and 0.045 mmol of AgNO3 were used.

Synthesis of NiFe LDH
The following raw materials: Ni(NO3)2·6H2O (58.29 mg), Fe(NO3)3·9H2O (58.27 mg), urea (51.91 mg) were dissolved in 25 mL of water and stirred for 20 min. The obtained solution and the nickel foam with a size of about 2.8 cm × 3 cm were then transferred to a 50 mL Teflon-lined stainless steel autoclave and hydrothermally heated at 120 °C for 12 h. After the autoclave cooled to room temperature (25 °C), the nickel foam was taken out, and the surface of the nickel foam was rinsed alternately with ethanol and water, and vacuum dried to obtain NiFe-LDH.

Characterization techniques
The crystal structure of the catalysts was analyzed using X-ray diffraction (Bruker D8 Advance, Cu Kα radiation). Morphological details were examined through SEM (JEOL JSM-7600F) and TEM (JEOL JEM-2100F). The electrode surface was dried using dust-free paper to remove residual moisture before XRD and SEM measurements. XPS analysis was conducted with a Thermo ESCALAB 250XI PHI-5000 system to explore the chemical valence states of the samples. Prior to XPS measurements, the electrode after ECR was thoroughly rinsed with deionized water to remove K2SO4 and carbonates, then dried in an oven at 60 °C. A PerkinElmer Avio 500 Inductively Coupled Plasma-Optical Emission Spectrometer (ICP-OES) was applied to quantify the elemental content. The potassium content was calculated relative to the total mass of the electrode, including the carbon paper, catalysts, and formed K2CO3 and KHCO3. Raman spectra were measured on Renishaw Raman microscope excited with a 532 nm laser. Before the Raman test, the GDE was removed from the electrolytic cell and used dust-free paper to absorb the electrolyte on the electrode surface instead of rinsing the surface to avoid removing the water-soluble (bi)carbonates. In situ shell-isolated nanoparticle-enhanced Raman spectroscopy (SHINERS) analysis was conducted using a Renishaw InVia Qontor confocal Raman microscope with a laser wavelength of 785 nm. A custom-designed flow cell was employed to collect Raman signals from a gas diffusion electrode loaded with the catalyst. The electrolyte consisted of 0.5 M K2SO4 (pH=1.8), with CO2 introduced to the back side of the GDE at a flow rate of 20 sccm. An Ag/AgCl electrode and a carbon rod were used as the reference and counter electrodes, respectively. An Olympus LUMPLANFLN 40× water dipping objective was immersed directly into the electrolyte to acquire the Raman spectra. For the in-situ ATR-SEIRAS (Attenuated total reflectance-surface-enhanced infrared absorption spectroscopy) experiments, Bi0.5-Ag0.5 and Bi catalysts deposited on Au-coated Si prisms were employed as the working electrodes, while an Ag/AgCl electrode and a Pt foil served as the reference and counter electrodes, respectively. Potentiodynamic ATR-SEIRAS spectra were recorded in 0.5 M KHCO3 electrolyte saturated with CO2. During the entire measurement, CO2 was continuously purged into the electrochemical cell at 20 sccm to maintain electrolyte saturation and enhance mass transport. Infrared signals were collected using a Thermo Fisher IS50 spectrometer equipped with an MCT-A detector, with unpolarized IR radiation incident at 60° and a spectral resolution of 8 cm−1. The spectra are presented in absorbance mode as −log(I/I0), where I and I0 denote the intensities of the reflected radiation from the sample and the reference, respectively.

Electrochemical measurements
For ECR//CEM//OER, the electrochemical measurement was conducted in a flow cell with electrode areas of 1 cm × 1 cm (the electrolyte flow layer thickness: 5 mm for anolyte flow and 5 mm for catholyte flow) separated by a Nafion 117 membrane using a CHI760E electrochemical workstation. 5 mg commercial 40 wt% Pt/C catalysts or 10 mg BiAg catalysts were dispersed in 960 µL isopropanol alcohol and 40 µL Nafion solution. The resulting ink was ultrasonically to achieve a uniform suspension, then air-brushed onto a carbon paper to obtain a catalyst loading of 1 mg cm−2. The three-electrode system includes 40 wt% Pt/C, Ag/AgCl and BiAg catalysts, serving as counter electrode, reference electrode and working electrode, respectively. The reference electrode was calibrated using a reversible hydrogen electrode in H2-saturated 0.5 M H2SO4 electrolyte (pH=0), with Pt wire and graphite rod as working electrode and counter electrode, respectively. The electrochemical impedance spectroscopy measurement was performed at open circuit potential over the frequency range from 100,000 to 0.1 Hz with an amplitude of 5 mV. Typical solution resistance for 0.5 M K2SO4 (pH=1.8) was measured to be 7.6 Ω ± 0.1 Ω. 100% iR compensation was applied for all potentials. The pH of the solutions was measured using a Accumet AB150 benchtop pH meter. The ECR performance of Bi0.5-Ag0.5 was evaluated in acidic, neutral, and alkaline solution using 0.5 M K2SO4 + H2SO4 (pH=1.8 ± 0.02), 0.5 M K2SO4 (pH=7.43 ± 0.08), and 1 M KOH (pH=14 ± 0.05) as the electrolytes, respectively. Fresh electrolytes were prepared daily and stored in sealed polypropylene bottles for use on the same day. The potentials were calibrated to RHE scale using ERHE = EAg/AgCl + 0.059 × pH + 0.197. During the tests, CO2 gas was delivered into the cathodic compartment at a rate of 20 sccm (controlled by a Sevenstar mass flow controller). The chronopotentiometry tests were carried out at current densities of −100, −200, −300, and −400 mA cm−2 to compare the ECR performance of different catalysts. The gas products were analyzed with gas chromatography (Shimadzu 2014) during the reaction and liquid products collected afterward for Nuclear Magnetic Resonance (a Bruker 500 MHz spectrometer (AV500 NMR)) analysis. To evaluate the electrochemical surface area (ECSA) of different samples, a 0.5 M K2SO4 (pH=1.8) electrolyte saturated with Ar was used with a flow rate of 20 mL min−1 in a three-electrode system. Cyclic voltammetry was conducted at scan rates of 10, 20, 30, 40, and 50 mV s−1 to determine electrochemical double-layer capacitance (Cdl).

The 8 h chronopotentiometry tests for all-type electrolyzers were performed at a current density of 100 mA cm−2 in 0.5 M K2SO4 (pH adjusted to 1.8 by 0.5 M H2SO4) in a two-electrode system with electrode areas of 1 cm × 1 cm (the total electrolyte flow layer thickness: 10 mm). CO2 was delivered into the cathodic compartment at a rate of 20 sccm (controlled by a Sevenstar mass flow controller). The back side of the GDE was pre-treated with a 10 wt% PTFE solution to mitigate potential electrolyte flooding. Specifically, 4 mL of PTFE solution was air-sprayed onto a 10 cm2 carbon paper, followed by calcination at 330 °C for 30 minutes in a muffle furnace with a heating rate of 5 °C/min. Bi0.5-Ag0.5 and 40 wt% commercial Pt/C sprayed onto the above PTFE treated gas diffusion layer as cathode and anode, respectively. Specifically, for both ECR//CEM//OER and ECR//CEM//HOR, a Nafion 117 proton-exchange membrane was employed to separate the anolyte and catholyte. The anolyte and catholyte volume were both 50 mL with a flow rate of 20 mL min−1 for the 8 h stability test at 100 mA cm−2. For ECR//CEM//HOR and ECR//HOR, H2 gases were delivered into the anodic compartment at a rate of 10 sccm. For ECR//OER and ECR//HOR, the electrolyte volume was 50 mL with a flow rate of 20 mL min−1 for the 8 h stability test at 100 mA cm−2.

For the 8 h stability test conducted in the two-electrode ECR//CEM//HOR electrolyzer with electrode areas of 0.5 cm × 2 cm (with a total electrolyte flow layer thickness of 2 mm), two different catholytes were used: 0.5 M K2SO4 and 1 M KOH, each with a volume of 100 mL. For the 8 h stability test in the two-electrode ECR//CEM//OER electrolyzer with electrode areas of 0.5 cm × 2 cm, the catholyte was 0.5 M K2SO4 (pH adjusted to 1.8 by 0.5 M H2SO4). During the 8 h electrolysis, the pH value of the catholyte was recorded every hour using a Accumet AB150 benchtop pH meter. At the same intervals, the catholyte were collected for liquid-phase product analysis via NMR spectroscopy. Simultaneously, the outlet gas stream from the cathode was analyzed by GC to quantify the gas-phase products.

The HOR polarization curves under different H2 flow rates (controlled by a Sevenstar mass flow controller) and H2/Ar ratios were measured in a three-electrode system with electrode areas of 1 cm × 1 cm. Pt/C, Ag/AgCl, and Bi0.5-Ag0.5 were used as the working electrode, reference electrode, and counter electrode, respectively. Polarization curves were performed from 0 to 1.6 V (V vs. RHE) at a scan rate of 20 mV/s. The electrolyte was 0.5 M K2SO4 (pH=1.8), and the electrolyte flow rate was maintained at 8 mL/min. Typical solution resistance in 0.5 M K2SO4 (pH=1.8) was measured to be 7.6 Ω ± 0.1 Ω. 100% iR compensation was applied. The ECR performance was evaluated under different H2 flow rates (5, 10, 15, and 20 sccm) using a two-electrode configuration. Moreover, polarization curves were conducted in a three-electrode system under different H2/Ar gas, including pure H2, 80% H2 + 20% Ar, 50% H2 + 50% Ar, and 30% H2 + 70% Ar. In addition, ECR performance was evaluated in a two-electrode system at a current density of 100 mA cm−2 under the same gas conditions.

For the long-term stability tests of the ECR//HOR electrolyzer, electrochemical measurements were performed in a two-electrode configuration using flow cells with electrode areas of either 0.5 cm × 2 cm or 2 cm × 2 cm. The corresponding electrolyte flow layer thicknesses were 2 mm and 3 mm, respectively. Electrolyte volumes of 500 mL (1 cm2 cell) and 1000 mL (4 cm2 cell) were circulated at flow rates of 8 mL min−1 and 5 mL min−1, respectively, and refreshed every 24 hours to maintain stability. In both setups, the CO2 feed was humidified by bubbling through HCl solutions (0.05 M for the 1 cm2 cell) and 20 sccm and 30 sccm for the 1 cm2 and 4 cm2 cells, respectively.

We conducted CO stripping experiments on the commercial Pt/C gas diffusion electrode (GDE) after the long term stability test. After operation in the flow cell, the Pt/C GDE was carefully transferred to a conventional three-electrode cell using 0.5 M K2SO4 (pH adjusted to 1.8 by 0.5 M H2SO4) as the electrolyte, with a saturated calomel electrode as the reference electrode and a graphite rod as the counter electrode. Cyclic voltammetry was first performed by scanning from 0.05 V to 1.55 V vs. RHE at 20 mV s−1 to detect any CO oxidation peak. Subsequently, high-purity CO gas was bubbled through the electrolyte for 20 min while holding the working electrode potential at 0.05 V vs. RHE to ensure sufficient CO adsorption. The cell was then flushed with fresh N2-purged electrolyte for 10 min to remove residual CO. Finally, CO stripping was carried out by CV scanning twice from 0.05 V to 1.55 V vs. RHE at a scan rate of 20 mV s−1.

ηH+ is calculated according to the following Eq. (3):

ηH+ = Transfered H+ / Passedelectron × 100% = ([H+]final + [OH−]generated) / (Current × Time / F) × 100% (3)

For MOR tests in the three-electrode system, Hg/HgO, Pt foil and NiFe LDH were used as reference electrode, counter electrode and working electrode, respectively. All the potentials obtained in three-electrode system were calibrated to RHE scale using ERHE = EHg/HgO + 0.059 × pH + 0.098. For the membrane electrode assembly, NiFe LDH and 40 wt% Pt/C were used as anode and cathode, respectively. An ink-cast method was used to prepare the cathode electrode. Specifically, 5 mg commercial 40 wt% Pt/C was dispersed in 980 µL isopropanol alcohol and 20 µL Nafion mixing solution and ultrasonically dispersed evenly, followed by drip coating on a carbon paper with a loading mass of 1 mg cm−2. NiFe LDH electrode was pressed under 1.4 MPa at 80 °C for 2 min before using as the anode. The anode and cathode chambers were separated by the anion exchange membrane (PiperION® self-supporting AEMs, 40 microns). The volume of anode and catholyte are both 150 mL. 1 M KOH + 1 M MeOH and 1 M KOH were used as anolyte and catholyte with the flow rate of 8 mL min−1, respectively. Typical solution resistance of 1 M KOH was measured to be 2.9 Ω ± 0.1 Ω. 100% iR compensation was applied for all potentials. Chronopotentiometry measurements were performed at 100 mA cm−2.

For the integrated ECR//HOR and HER//AEM//MOR electrolyzer, electrolysis was performed at a current density of 100 mA cm−2 at the HER//AEM//MOR electrolyzer for 20 min to produce enough H2 and purge air from the catholyte and the electrolyte container, followed by supplying H2 into the ECR//HOR electrolyzer. For ECR//HOR, the electrolyte is 100 mL. For HER//AEM//MOR electrolyzer, the anolyte is 25 mL. All the electrolyte flow rates were set as 8 mL min−1.

Product analysis
The gas products and liquid products were analyzed with GC (Shimadzu 2014) and NMR, respectively. The liquid products were analyzed through 1H NMR spectroscopy on a Bruker 500 MHz spectrometer (AV500 NMR). Prior to analysis, the post-reaction electrolyte was neutralized with KOH and diluted with water. The 1H spectra were measured by using dimethyl sulfoxide as internal standard and the water peak was suppressed during the test. Specifically, 100 µL of the electrolyte was mixed with 100 µL of KOH, 400 µL of H2O, and 100 µL of D2O for the NMR test.

The FE of gas products are calculated as below:

FE = j_i / j_total × 100% = (x_i × v × n_i × F × p_o) / (R × T × j_total × electrode area) × 100% (4)

where j_total and j_i are total current density and and partial current density, respectively. x_i is the volume fraction of a given product determined by online GC, based on calibration curves established using a standard gas mixture. v represents the gas flow rate, fixed at 20 sccm; n_i is the number of electrons transferred per molecule of the product; p_o = 101.3 kPa; F is the Faraday constant (96485 C mol−1); T = 298 K and R denotes the gas constant.

The FE of liquid products are calculated as follows:

FE = (e × F × n) / Q = (e × F × n) / (I × t) (5)

Where n represents the moles of electrons required to generate the products, F denotes the Faraday constant; e is the number of electrons transferred; I is the applied current and t refers for the electrolysis time.

COMSOL simulation
Mass balance and charge conservation in the electrolytes and Nafion membrane domains were estimated. Nernst-Planck equation with acid-base balance relations was used to describe the mass balance and the charge conservation in the electrolytes and Nafion membrane domains.

In aqueous electrolyte domains:

∂Ci / ∂t = ∇(Di ∇Ci + zi Di F/RT Ci ∇φ) (6)

−∇(ε0 εr ∇φ) = F Σ zi Ci (7)

In Nafion membrane domains, there are charge neutrality with a constant concentration of −SO3− negative charges (c_SO3 = 1 mol/L) and a constant proton selectivity:

∂Ci / ∂t = ε ∇(Di ∇Ci + zi Di F/RT Ci ∇φ) / τ (8)

−∇(ε0 εr ∇φ) = F(Σ zi Ci − c_SO3) (9)

k = H+ / c_SO3 (10)

To better explore the effect of a single variable on mass transfer behavior, there are a constant H-ion concentration (0 mol/L) and a constant potential (0 V) at the cathode-catholyte interface, and there are a constant current density (I=100 mA cm−2) and a constant potential (1 V) at the anode-anolyte interface. Mobile ions included H+, K+, and SO42−. The Nafion membrane carried a fixed negative charge (c_SO3 = 1 mol/L), and ion selectivity was described by an effective H+ transport number (αH+). The baseline parameters were an electrolyte composition of 0.5 M K2SO4 + 0.01 M H2SO4, a membrane thickness of 200 µm, and a proton selectivity of 0.8. Individual parameters were then varied to assess their specific effects.

All finite-element simulations were conducted in COMSOL Multiphysics. Unless otherwise specified, all other parameter conditions were set to be consistent with the electrochemical measurement conditions as shown in Supplementary Table 1.

Peer review rounds
[{"round": 1, "reviewer_comments": {"reviewer_1": "The manuscript titled \"Membrane-free CO2 hydrogenation electrolyzer for salt precipitation management in acidic electrochemical CO2 reduction\" addresses a practical limitation in acidic CO2 electroreduction (ECR) systems — salt precipitation due to proton transport inefficiency. By proposing a membrane-free electrolyzer coupled with hydrogen oxidation reaction (HOR), the authors aim to stabilize catholyte pH, suppress carbonate formation, and improve system stability. The study includes thoughtful electrochemical evaluations, a techno-economic analysis (TEA), and mechanistic insights into ion transport. While, the novelty is limited, particularly regarding the electrolyzer configuration. The proposed strategies echo existing literature without providing sufficient mechanistic differentiation or rigorous theoretical validation. The role of ηH+ is discussed as a central concept, yet it remains empirically defined without being quantitatively supported by theoretical models. Therefore, this manuscript is not suitable for publication in Nature Communications currently.\n1. The membrane-free + HOR configuration has been explored previously (e.g., Joule, Volume 6, Issue 12, 2745-2761). The authors need to clearly explain the difference between their work and previous ones.\n2. The definition of ηH+ is operational and lacks a rigorous basis. No modeling (e.g., transport simulations) is presented to support how membrane properties or electrolyte composition influence proton transfer.\n3. The authors neglect possible effects of hydrogen concentration variation, gas crossover, or local pH shifts at the anode during long-term operation.\n4. Is the Bi oxidation state truly preserved under acidic bias?\n5. The rationale for using Bi0.5Ag0.5 is mainly based on conductivity enhancement, which is not novel. The role of Bi-Ag interaction in intermediate stabilization or CO2 binding needs to be better elaborated.\n6. Supplying pure H2 to the anode may not be economically viable for large-scale systems. Authors should at least discuss alternative hydrogen sources (e.g., internal HER coupling) or system integration strategies.\n7. To highlight the advantage of the membrane-free acidic system, it would be informative to include direct performance comparisons (FE, energy efficiency, voltage) with neutral or alkaline membrane-based ECR systems using the same Bi–Ag catalyst.\n8. Several key characterizations, including catalyst structure, long-term stability data, and electrolyte compositions, are buried in the supplementary section. Some of these should be promoted to the main text to ensure scientific transparency.", "reviewer_2": "This manuscript reported a membrane-free reactor for CO2 reduction in an acidic electrolyte. Compared with membrane-based electrolyzers, the membrane-free design maintains a stable electrolyte pH, which is crucial for acidic CO2 reduction. To address challenges associated with this design, they utilized electrochemical hydrogen oxidation to replace the oxygen evolution reaction, thereby suppressing formic acid overoxidation at the anode. Additionally, they proposed coupling methanol oxidation with the hydrogen evolution reaction to reduce H2 costs. As a result, the reactor exhibited high selectivity, single-pass efficiency, and stability. More impressively, they emphasized the often-overlooked significance of proton transfer efficiency across the membrane, which greatly affects the pH of the catholyte. This is important for accurately evaluate acidic CO2 reduction performance and maintain long-term stability. Therefore, this manuscript could be accepted for publication in Nature Communications after addressing the following questions.\n1. The authors used the same pH solution as both the anolyte and catholyte. However, some acidic CO2 reduction studies use a pure acid solution as the anolyte. What is the influence of the anolyte composition?\n2. The authors developed BiAg catalyst for CO2-to-formic acid conversion. How does its performance compare with previously reported catalysts? Additionally, how does it’s performance in neutral and alkaline electrolytes?\n3. In the membrane-based electrolyzer design, why does the catholyte pH stabilize at ~7-8 after several hours of testing?\n4. The authors reported a 50-hour acidic stability. How does this result compare with current literature values?\n5. The TEA analysis did not clearly explain the calculation of the liquid product separation cost, which needs to be refined for greater clarity. Furthermore, the liquid product separation cost should vary among different electrolyzers due to differences in catholyte usage. The HER//AEM//MOR system should also account for the separation cost. The transformation cost from formate to formic acid should also be considered in the HER//AEM//MOR system.\n6. Some statements in the SI are inaccurate, such as electrocatalytic ECR activity. Please carefully check the statements.", "reviewer_3": "In this study, the authors propose a membrane-free electrolyzer to enhance the stability and economic viability of electrocatalytic CO2 reduction in acidic systems. Experimental results demonstrate that the membrane-free electrolyzer significantly improves operational stability from 8 hours to 50 hours. To further optimize the system's economic efficiency, the authors employ hydrogen oxidation reaction (HOR) to replace the conventional oxygen evolution reaction (OER) at the anode and introduce a tandem reaction strategy (HOR-CEM-MOR) to reduce overall operational costs. Upon careful evaluation, the reviewer considers the current study presents critical shortcomings in practical applicability and data accuracy that undermine the validity of its conclusions. Specific concerns are outlined below:\n1. The experimental results report a dilute formic acid concentration of 0.1 M was produced in CO2 reduction system. However, the techno-economic analysis (TEA) lacks rigor, as the authors fail to account for the energy and costs required to concentrate and purify the product to ≥95% purity, which is essential for accurate economic evaluation.\n2. To fairly highlight the advantages of the membrane-free strategy, comparative analyses under equivalent conditions are imperative. Notably, the parameters for the CRR//CEM//OER process were derived from a 10-mm flow channel plate, whereas those for CRR//HOR and CRR/CEM/MOR were based on a 2-mm plate. This discrepancy invalidates direct economic comparisons.\n3. As the authors emphasize that acidic systems avoid the acidification and separation costs associated with alkaline-derived formate. However, the CRR//CEM//MOR configuration employs an alkaline electrolyte, this introduces additional formic acid purification expenses, which are not considered in the current TEA.\n4. Carbonate formation typically limits CO2 reduction stability to <200 h due to pore-blocking effects in CO2 transport channels. Despite adopting a membrane-free system, the reported stability (50 h) remains insufficient to demonstrate practical superiority. Furthermore, this limited stability raises concerns about operational feasibility in real-world applications.\n5. The reviewer questions whether reactant crossover from the anode to the cathode via the gas diffusion layer (GDL) in the membrane-free system would increase gas separation costs, a factor omitted in the analysis.\n6. The manuscript lacks a comprehensive list of parameters (e.g., energy inputs, material costs, separation efficiencies) used in the TEA, hindering reproducibility and critical assessment.\n7. The electrocatalytic reaction area is not explicitly stated. Detailed experimental information, including electrode dimensions and active surface area, must be provided.\n8. The currency unit \"dollars\" should be standardized as \"USD\" throughout the manuscript for clarity and consistency.", "reviewer_4": "This article by Da et al. showed a detailed study on using a membrane-free electrolyzer design to manage pH at the cathode and hence mitigate salt precipitation during CO2 electrolysis, which is a major challenge in the field. They systematically compared the effect of pH stabilization with different electrode pairing strategies (ECR/CEM/OER, ECR/CEM/HOR, ECR/OER, and ECR/HOR), identifying the membraneless ECR/HOR as the most promising configuration in terms of formic acid faradaic efficiency, proton transfer efficiency, pH management, and cell voltage, etc. They also demonstrated a cascade of ECR/HOR and ECR/AEM/MOR electrolyzers that show promise for practical application based on the TEA. Although the concept of membrane-free CER/HOR electrolyzers is not new (see Fu et al., Science 379, 707–712 (2023); Pang et al. Joule 6, 12, 2745 - 2761; and Lazouski et al. Nat Catal 3, 463–469 (2020), none of which were cited), the authors added value by comparing the different cell configurations and extending to a tandem scheme while achieving appreciable performance in terms of current density, Faradaic efficiency, and single pass conversion efficiency. The study presented here should be of interest to the electrochemistry and the broader CO2 utilization communities. However, the following concerns need to be addressed before the article is suitable for publication.\n1. The TEA looks promising and can be helpful for projecting the competitive advantage of the new technology, but I have some questions and concerns. (i) A more detailed written description of the assumptions for “the liquid product separation” is required. Currently, the method(s) of separation are not described, key assumptions about the concentration of the solutions entering and leaving the separation unit are absent, and there are no citations for the Capex and operating costs. For the tandem process involving formate production from MOR, was the liquid separation cost for the MOR-derived formate accounted for as well? Salty, alkaline formate/formic acid will have very little value. (ii) The TEA should be carried out as a discounted cash flow analysis using a typical discount rate (e.g. 10%). It seems that this was not done. (iii) Was the electrolyzer Capex of $919/m^2, derived from a value of $250 $/kW (which is very low), applied to both the water electrolyzer and the CO2 electrolyzer? This should be clearly stated. It’s also not clear if the 250 $/kW is for a PEM or liquid alkaline electrolyzer. (iv) It would be very helpful to briefly state some of the key assumptions for the TEA in the main article (e.g., cost of electrolyzer, Capex, method of separation, price of CO2, price of MeOH) so that the reader can better understand the baseline scenario being presented. Optional: a sensitivity analysis would be helpful.\n2. Regarding Figure 3, although the ECR/HOR configuration and the 2 mm gap significantly lowered the cell voltage, 1.7 V is still a relatively high cell voltage, especially if ECR (-0.2 V vs. SHE) and HOR （0 V vs. SHE) are paired. Can the authors break down the overpotential losses? Also, with the ECR/OER and ECR/CEM/OER configurations, the cell voltages are between 5 and 6 V. Did the authors observe any degradation of the carbon paper substrate as carbon is prone to oxidation? For the ECR/HOR configuration without a membrane, why is the cell voltage so high (~6 V)? What is the origin of the resistance?\n3. How scalable is the membraneless electrolyzer design? It seems that all results for the CO2 electrolyzer were obtained using very small (1 cm^2) electrodes. Did the authors attempt any measurements at a larger scale to see how/if performance changes?\n4. How stable can the electrolyzer system be beyond 50 hours, which is a relatively short time for a durability test for electrolyzer technology?\n5. Since CO was observed as shown in Figure 2, did the authors observe any CO poisoning effect with the Pt anode? If so, how significant is it, especially during long-term operation?\n6. Regarding Supplementary Figure 21, the bicarbonate peak should be overlaid with and compared to the other configurations as shown in Supplementary Figure 19, especially since the y-axis does not indicate the scale.\nMinor comments:\nThe authors should also list all the assumptions made during the simulation of the scenarios as shown in Figure 1 b-c.\nThe authors should define MOR (methanol oxidation reaction) when they first introduced it.", "reviewer_5": "I co-reviewed this manuscript with one of the reviewers who provided the listed reports. This is part of the Nature Communications initiative to facilitate training in peer review and to provide appropriate recognition for Early Career Researchers who co-review manuscripts."}, "author_response": "We sincerely thank the reviewer for their positive comments and constructive suggestions. We have carried out additional CO electroreduction tests under neutral and alkaline conditions to demonstrate the catalyst’s performance, optimized the catalyst reaction conditions to extend the long-term stability to 208 hours, and included an updated techno-economic analysis with revised formic acid separation costs. We have also added a detailed comparison with recent membrane-free and membrane-based electrolyzers to highlight the advantages of our system in selectivity and stability. These new results and discussions have been incorporated into the revised manuscript. We believe these improvements comprehensively address the reviewer’s concerns and further strengthen the impact and practical significance of our work.\n1. The authors used the same pH solution as both the anolyte and catholyte. However, some acidic CO reduction studies use a pure acid solution as the anolyte. What is the influence of the anolyte composition?\nResponse:\nWe sincerely thank the reviewer for the insightful comment. The composition of the anolyte and catholyte, including the anolyte species (pH, with or without alkaline cations), concentration, as well as the concentration gradient between the anolyte and catholyte, significantly influence J i, diffusion. In other words, the anolyte can directly affect the catholyte composition and pH. Although the anolyte is separated from the cathode by a membrane in previous membrane-based design, it still impacts the cathodic CO reduction reaction through proton crossover and pH drift. A pure acidic anolyte may enhance H+ flux to the cathode, helping maintain a low catholyte pH; however, excessive proton transport can result in ηH+ exceeding 100%, leading to continuous pH decline at the cathode, which in turn promotes HER and suppresses CO reduction selectivity. In contrast, our membrane-free design enables strict control of ηH+ at 100%, maintaining a stable electrolyte pH throughout operation.\n2. The authors developed BiAg catalyst for CO2-to-formic acid conversion. How does its performance compare with previously reported catalysts? Additionally, how does it’s performance in neutral and alkaline electrolytes?\nResponse:\nWe sincerely thank the reviewer for the valuable question regarding the performance comparison of our Bi-Ag catalyst. To address this point, we conducted additional experiments using a three-electrode flow cell system to evaluate the intrinsic CO2 reduction performance of the Bi0.5Ag0.5 catalyst in both neutral (0.5 M K2SO4) and alkaline (1 M KOH) electrolytes with an ECR//CEM//OER electrolyzer. As shown in Supplementary Fig. 12, the catalyst exhibited excellent ECR selectivity across a wide range of current densities (100−400 mA cm−2). In both electrolytes, the formic acid FEs remained above 90%, demonstrating the versatility of the Bi0.5Ag0.5 catalyst in different pH environments.\nIn addition, our Bi0.5Ag0.5 catalyst delivers a Faradaic efficiency for formic acid over 90% at 100 mA cm−2 in acidic electrolytes, while maintaining stability for 208 hours. Several previously reported systems have demonstrated CO2-to-formic acid electrolysis at higher current densities (Supplementary Table 4); however, their operation lifetimes under acidic conditions are typically limited to several tens of hours due to catalyst deactivation or system instability. In contrast, the extended durability of our catalyst under acidic conditions at industrially relevant current densities underscores its potential for practical large-scale application.\nSupplementary Table 4. Comparison of the performance of Bi0.5Ag0.5 with catalysts reported in representative literature.\nCatalyst Electrod e area (cm2) Product Catholyte pH Current density (mA cm−2) Time (h) Ref.\nSn SAC \\ HCOOH 0.5 M K2SO4 + H2SO4 3 150 18 32\nSn(S)-H 1 HCOOH 0.5 M K2SO4 + H2SO4 3 400 13.5 33\nBi nanorods 1 HCOOH 0.5 M K2SO4 + H2SO4 2 500 50 34\nCu Sn 4 HCOOH 3 M KCl + H2SO4 1 100 130 35\nSnBi 2.89 HCOOH 3 M KCl + H2SO4 1 100 125 36\nNU-1000-Sn 0.2 HCOOH 3 M KCl + H2SO4 1.6 260 15 37\nBi-MOF-TS 0.25 HCOOH 0.1 M K2SO4 + H2SO4 ± 0.2 1.5 500 60 38\nr-Pb 1 HCOOH 0.5 M K2SO4 + H2SO4 1 600 5200 10\nBi0.5Ag0.5 1 HCOOH 0.5 M K2SO4 + H2SO4 1.8 100 208 This work\nBi0.5Ag0.5 4 HCOOH 0.5 M K2SO4 + H2SO4 1.8 100 48 This work\nSupplementary Fig. 12 | Products distribution of Bi0.5Ag0.5 at various applied current densities in different electrolytes. a, Products distribution of Bi0.5Ag0.5 at various applied current densities in ECR//CEM//OER electrolyzer in 1 M KOH. b, Products distribution of Bi0.5Ag0.5 at various applied current densities in ECR//CEM//OER electrolyzer in 0.5 M K2SO4.\nWe supplemented the information in the manuscript and experimental details:\nMain text, Page 7\nThe ECR performance of the Bi0.5Ag0.5 catalyst was also assessed in neutral and alkaline electrolytes, where formate Faradaic efficiencies consistently exceeded 90% (Supplementary Fig. 12).\nMain text, Page 17\nThe ECR performance of Bi0.5Ag0.5 was evaluated in acidic, neutral, and alkaline solution using 0.5 M K2SO4 + H2SO4, 0.5 M K2SO4, and 1 M KOH as the electrolytes, respectively.\n3. In the membrane-based electrolyzer design, why does the catholyte pH stabilize at ~7-8 after several hours of testing?\nResponse:\nWe thank the reviewer for this insightful question. In the membrane-based electrolyzer, the catholyte pH initially increases during electrolysis due to the consumption of protons and generation of hydroxide ions at the cathode during CO2 reduction. However, after several hours, the catholyte pH stabilizes around 7-8 due to the formation of a bicarbonate/carbonate buffer system. As CO2 dissolves in the aqueous electrolyte, it reacts with OH− generated at the cathode to form HCO3−, thereby establishing a chemical buffer that resists further pH change. Additionally, in the case of CEM-based systems, the migration of H+ from the anode to the cathode further contributes to moderating the catholyte pH, preventing it from rising to strongly alkaline values. The combined effects of buffering reactions and membrane ion transport explain the observed pH stabilization.\n4. The authors reported a 50-hour acidic stability. How does this result compare with current literature values?\nResponse:\nWe appreciate the reviewer’s comment regarding the comparison with literature-reported stability data. In our revised study, we optimized the electrolysis conditions to further enhance system durability under acidic environments. Specifically, to suppress carbonate precipitation and maintain stable operation, the CO2 feed was humidified by passing through a glass bubbler containing 0.05 M HCl prior to entering the cathode chamber of the ECR//HOR electrolyzer (Science 2025, 388, 6752, eadr3834). The introduction of volatile acids helps convert carbonate into soluble potassium salts, thereby mitigating pore blockage.\nElectrolysis tests were conducted using gas diffusion electrodes (GDEs) with exposed areas of 1 cm2 (electrolyzer dimension: 0.5 cm × 2 cm) and 4 cm2 (electrolyzer dimension: 2 cm × 2 cm). Catalysts were uniformly spray-coated to ensure complete active area coverage and direct exposure to both gas and electrolyte phases during operation. Under these optimized conditions, the system achieved continuous operation for 208 hours at a current density of 100 mA cm−2, with the Faradaic efficiency for HCOOH stably maintained at ~ 90% in the 1 cm2 cell. This significantly exceeds the previously reported 50-hour stability and underscores the robustness of our membrane-free design. Additionally, in a 4 cm2 electrolyzer, similar performance (~ 90% FE for HCOOH) was sustained for 48 hours.\nWe have included a detailed comparison of our stability data with representative state-of-the-art acidic ECR systems in the revised manuscript (see Supplementary Table 4). It shows that our system delivers comparable or superior long-term stability and product selectivity relative to most state-of-the-art reports.\nWe have added the relevant discussion in the main text on page 19 and Supplementary Table 4.\nMain text, Page 19\nFor the long-term stability tests of the ECR//HOR electrolyzer, electrochemical measurements were performed in a two-electrode configuration using flow cells with electrode areas of either 0.5 cm × 2 cm or 2 cm × 2 cm. The corresponding electrolyte flow layer thicknesses were 2 mm and 3 mm, respectively. Electrolyte volumes of 500 mL (1 cm2 cell) and 1000 mL (4 cm2 cell) were circulated at flow rates of 8 mL min−1 and 5 mL min−1, respectively, and refreshed every 24 hours to maintain stability. In both setups, the CO2 feed was humidified by bubbling through HCl solutions (0.05 M for the 1 cm2 cell and 0.2 M for the 4 cm2 cell) prior to entering the cathode compartment8. CO2 flow rates were set at 20 sccm and 30 sccm for the smaller and larger cells, respectively.\nSupplementary Table 4. Comparison of the performance of Bi0.5Ag0.5 with catalysts reported in representative literature.\nCatalyst Electrod e area (cm2) Product Catholyte pH Current density (mA cm−2) Time (h) Ref.\nSn SAC \\ HCOOH 0.5 M K2SO4 + H2SO4 3 150 18 32\nSn(S)-H 1 HCOOH 0.5 M K2SO4 + H2SO4 3 400 13.5 33\nBi nanorods 1 HCOOH 0.5 M K2SO4 + H2SO4 2 500 50 34\nCu Sn 4 HCOOH 3 M KCl + H2SO4 1 100 130 35\nSnBi 2.89 HCOOH 3 M KCl + H2SO4 1 100 125 36\nNU-1000-Sn 0.2 HCOOH 3 M KCl + H2SO4 1.6 260 15 37\nBi-MOF-TS 0.25 HCOOH 0.1 M K2SO4 + H2SO4 ± 0.2 1.5 500 60 38\nr-Pb 1 HCOOH 0.5 M K2SO4 + H2SO4 1 600 5200 10\nBi0.5Ag0.5 1 HCOOH 0.5 M K2SO4 + H2SO4 1.8 100 208 This work\nBi0.5Ag0.5 4 HCOOH 0.5 M K2SO4 + H2SO4 1.8 100 48 This work\nReferences:\n32. Sun, B. et al. Unveiling pH-dependent adsorption strength of *CO2- intermediate over high-density Sn single Atom catalyst for acidic CO2-to-HCOOH electroreduction. Angew. Chem. Int. Ed. 63, e202318874 (2024).\n33. Shen, H. et al. Acidic CO2-to-HCOOH electrolysis with industrial-level current on phase engineered tin sulfide. Nat. Commun. 14, 2843 (2023).\n34. Li-Ping Chi et al. Efficient and stable acidic CO2 electrolysis to formic acid by a reservoir structure design. Proc. Nat. Acad. Sci. 120, e2312876120 (2023).\n35. Yu, X. et al. Coverage enhancement accelerates acidic CO2 electrolysis at ampere-level current with high energy and carbon efficiencies. Nat. Commun. 15, 1711 (2024).\n36. Li, L., Liu, Z., Yu, X. & Zhong, M. Achieving high single-pass carbon conversion efficiencies in durable CO2 electroreduction in strong acids via electrode structure engineering. Angew. Chem. Int. Ed. 62, e202300226 (2023).\n37. Xue, H., Zhao, Z.H., Liao, P.Q. & Chen, X.M. \"Ship-in-a-bottle\" integration of ditin(IV) sites into a metal-organic framework for boosting electroreduction of CO2 in acidic electrolyte. J. Am. Chem. Soc. 145, 16978-16982 (2023).\n38. Chen, X. et al. Activating inert non-defect sites in Bi catalysts using tensile strain engineering for highly active CO2 electroreduction. Nat. Commun. 16, 1927 (2025).\n5. The TEA analysis did not clearly explain the calculation of the liquid product separation cost, which needs to be refined for greater clarity. Furthermore, the liquid product separation cost should vary among different electrolyzers due to differences in catholyte usage. The HER//AEM//MOR system should also account for the separation cost. The transformation cost from formate to formic acid should also be considered in the HER//AEM//MOR system.\nResponse:\nWe appreciate the reviewer’s constructive comments. We have revised the TEA analysis that consider the products separation cost in different electrolyzers (ECR//HOR and HER//AEM//MOR) using different separation methods in the revised Supplementary Information, also copied below:\nSupplementary Information, Page 51\na) Liquid separation cost in ECR//HOR electrolyzer\nFor the liquid product separation assumption, a pressurized extraction-distillation method was adopted based on a previously reported process design7. In this approach, the formic acid-containing electrolyte is first processed through an extraction distillation column, followed by solvent recovery in a second column. Through this two-step separation, formic acid is concentrated to over 95 wt%, meeting commercial purity standards. All process parameters and energy consumption estimates were referenced from the literature7. The estimated energy consumption for the two columns is approximately 1097.2 kWh and 50.1 kWh, respectively.\nAccordingly, the liquid separation operational cost can be calculated as\nLiquid Separation Operational Cost (USD/tonne) = (1097.22 kWh+50.11 kWh) / 89.7 kWh * 1000 * 0.03 = 383.7 USD/tonne FA\nLiquid product flow rate (m3/hr) = 100*10^6 g / (1.02 g/ml * 24 hr * 10%) = 40.85 (m3/hr)\nCathode distillation capital cost (USD) = Distillation ref cost (USD/tonne) * Liquid product flow rate (m3/hr) * 1000 (L/m3) * (scale factor) / (Distillation ref scale (L/min) * 60 min/hr) = 2,978,658.8 USD\nLiquid Separation capital cost (USD/tonne) = Cathode distillation capital cost (USD) * (Liquid product flow rate (m3/hr) * 1000 (L/m3) * (scale factor) / (Distillation ref scale (L/min) * 60 min/hr)) / (Plant capacity factor * 365 day/year * 100 tonne/day) = 6.13 USD/tonne FA\nTotal cost of liquid product (USD/tonne) = Liquid Separation capital cost (USD/tonne) + Liquid Separation Operational Cost (USD/tonne) = 389.8 (USD/tonne FA)\nSupplementary Information, Page 58\nb) Liquid separation cost in HER//AEM//MOR electrolyzer\nFor the alkaline HER//AEM//MOR system, the primary product is formate rather than formic acid. According to the study by Mahinder Ramdin21, the purification of formate involves electrodialysis followed by distillation. The capital expenditure (CAPEX) associated with the electrodialysis and distillation processes is estimated at 0.06 USD kg−1 to deliver at least 85 wt% formic acid, while the operating expenditures (OPEX) for electrodialysis and distillation are 0.13 USD kg−1 and 0.12 USD kg−1, respectively. The associated cost of liquid-phase separation is therefore estimated with reference to the separation cost calculations reported in that work.\nLiquid phase separation CAPEX = 50000 kg/day * 365 day * (0.06+0.06) USD/kg * 20 year = 43800000 USD\nLiquid phase separation OPEX = 50000 kg/day * 365 day * (0.12+0.13) = 4562500 USD\nTherefore, the total liquid phase separation is\nLiquid phase separation cost (USD/tonne) = (43800000 USD+4562500 USD*20 year) / (50 tonne/day * 365 day * 20 year) = 370 USD/tonne\n6. Some statements in the SI are inaccurate, such as electrocatalytic ECR activity. Please carefully check the statements.\nResponse:\nWe appreciate the reviewer’s careful reading. We have thoroughly checked the statements in the Supporting Information, especially those related to the electrocatalytic ECR activity, and have corrected to \"ECR performance\".\nSupplementary Fig. 8 | ECR performance of Bi xAg1-x. a, ECR performance of Bi. b, ECR activity of Bi0.3Ag0.7 alloy. c, ECR performance of Bi0.7Ag0.3 alloy. The HCOOH FEs of all the catalysts are ~ 90% over the current density rage of 100−400 mA cm−2."}, {"round": 2, "reviewer_comments": {"reviewer_1": "I have carefully reviewed the revised manuscript and the authors’ point-by-point responses to my initial comments. The authors have provided additional experiments, analyses, and discussions that partially address my concerns, including performance comparisons with membrane-based systems, qualitative discussion of ηH+ transport, and mechanistic insights into the Bi–Ag catalyst. These revisions improve the clarity and completeness of the work. However, several of my original concerns remain only partially resolved.\n1. While the Nernst–Planck equation is now introduced, no quantitative modeling or parameter measurement is provided. Please consider simplified transport simulations (e.g., COMSOL) to quantify the effects of membrane thickness, ion selectivity, and electrolyte composition on ηH+.\n2. The SHINERS data suggest a lower onset potential for *CO2- formation on Bi0.5–Ag0.5, but the mechanistic explanation remains focused on conductivity. Please elaborate on whether Ag modifies the adsorption energies of *CO2- or *OCHO, or alters the electronic structure of Bi.\n3. The proposed MOR–HER coupling is presented as a proof-of-concept, but lacks quantitative energy and cost analysis. Please compare the energy efficiency of external H2 supply versus internal H₂ generation (including MOR energy demand and CO2 byproduct management), and provide a preliminary TEA for industrial-scale operation.", "reviewer_2": "The authors have supplemented related experimental results, and have well addressed most of my comments. I recommend the publication of this manuscript.", "reviewer_3": "By reviewing the point-by-point response to the reviewers, the reviewer believes that the authors' citation and methodology regarding TEA still have significant flaws. For example, in the response to comment 1, Reference 7 does not involve the calculation of TEA, so the reviewer could not verify the accuracy of the authors' citation. (All process parameters and energy consumption estimates were referenced from the literature7. Reference 7. Du, Z.Y. et al. Promoting Water Activation via Molecular Engineering Enables Efficient Asymmetric C-C Coupling during CO2 Electroreduction. J. Am. Chem. Soc. 146, 32870-32879 (2024).)\nSimilarly, in Comment 3, I could not find the corresponding description—Reference 21 does not include an author named Mahinder Ramdin. (According to the study by Mahinder Ramdin21. 21. Bazmi, M., Gong, J., Jessen, K. & Tsotsis, T. Waste CO2 capture and utilization for methanol production via a novel membrane contactor reactor process: Techno-economic analysis (TEA), and comparison with other existing and emerging technologies. Chem. Eng. Process.: Process Intensif. 201 (2024).）\nTherefore, regarding these two issues about TEA, I believe the authors have not addressed my concerns because I cannot assess the accuracy of their response. The data sources provided by the authors are unreliable, so I do not recommend accepting this article for publication in this journal.\nFinally, I noticed that on page 12 of the manuscript, it mentions that the formic acid production concentration is 0.1 M. Thus, the reviewer suggests that the authors should base their TEA calculations for formic acid separation and concentration on this experimentally obtained value to ensure accurate computations.", "reviewer_4": "In the updated TEA calculations, I believe there is a significant error in the calculation of the Capex for the CO2 electrolyzer. Specifically, on page 50 of the response document where the authors shown their calculation of the area-normalized Capex based on the reported Capex of traditional PEM water electrolyzers (~ $560 USD/kW), they assume a current density of 0.17 A/cm2. This is incorrect. Conventional PEM electrolyzers should have a reference current density of closer to 1.7 A/cm2, not 0.17 A/cm2. As a result, the area-based Capex should be $16,184 USD/m^2, NOT $1,618 USD/m^2. As an aside, the authors should be careful with their dimensional analysis and more clearly spell out the conversion from kW to W and cm^2 to m^2.\nSupplementary Fig. 39 (sensitivity analysis)- I find it confusing to interpret this figure. I suggest that the authors re-draw it in the form of a more traditional “tornado plot” where the baseline, “high”, and “low” values of the dependent variable are all plotted on the same axis instead of two identical axes as was done here. Furthermore, I suggest the authors plot the levelized cost of FA production as the dependent variable being plotted rather than NPV. NPV is not very useful on its own since it is so dependent on plant size.\nRegarding scalability, it is encouraging that the new larger cell also demonstrated high FE towards FA, but the authors should also add overlaid polarization (current density vs. voltage) curves for both the 0.5 cm × 2 cm and 2 cm × 2 cm cells. This will be very helpful to see how the performance is affected by the increased electrode size.\nRegarding CO poisoning of the Pt anode catalyst: I do not agree with the authors' interpretation of their new results. Supplementary Fig. 28 clearly shows a CO stripping peak around +0.8 V vs. RHE; below this potential it is likely that Pt has a difficult time doing HOR (as is well known in H2 fuel cell literature). That means that the anode probably needs at least +0.8 V of extra overpotential to oxidize CO and be active towards HOR. i.e., at high enough voltage, the poison will be removed, but device efficiency suffers from the higher voltage. The lack of a CO stripping peak after the stability test doesn’t prove that CO is not crossing over, and it doesn’t mean that CO poisoning is not an issue. I would guess that the use of a more CO tolerant anode catalyst like PtRu would lead to significantly reduced cell voltage. Thus, I still disagree with the authors’ statement that there is no CO crossover in their cell in this line: “ These results demonstrate that the CO generated at the cathode did not crossover to the anode to cause CO poisoning under the tested flow cell conditions.” At the least, I feel like this statement is very misleading and should be changed.", "reviewer_5": "I co-reviewed this manuscript with one of the reviewers who provided the listed reports. This is part of the Nature Communications initiative to facilitate training in peer review and to provide appropriate recognition for Early Career Researchers who co-review manuscripts."}, "author_response": "Thanks for your valuable suggestions, which have significantly improved our work. Following your advice, we incorporated additional COMSOL simulations to examine the effects of electrolyte composition and membrane properties on proton transport and their impact on ηH+. For the catalytic mechanism, we employed in situ attenuated total reflection surface-enhanced infrared absorption spectroscopy (ATR-SEIRAS) to probe intermediate binding and used XPS to investigate electron transfer on Bi. We also expanded the discussion to provide deeper mechanistic insights into the Bi-Ag catalyst. Regarding the TEA, we revised the analysis using Aspen-based process simulations that account for H2 supply, byproduct management, and industrial-scale operation, enabling a more rigorous and comprehensive evaluation."}, {"round": 3, "reviewer_comments": {"reviewer_1": "The questions have been properly addressed. I believe this work is ready for publication now.", "reviewer_3": "I have carefully reviewed the authors' detailed content regarding the TEA results and appreciate their contributions in this area, which have enhanced the rigor of the scientific research. However, I believe the authors have not fully addressed the issues I raised. My primary concern lies in the cost consumption during the concentration process of the dilute formic acid stream obtained in this experiment. In their response to Comment 2, the authors cited content from Reference 9 and Reference 12 to support their analysis of the concentration cost of the formic acid product. However, the reviewer noted that in Reference 9, the original authors used a formic acid feed concentration of 10%wt, rather than the concentration of 0.1 mol/L ≈ 0.46%wt obtained in the authors' article. Therefore, the reviewer considers it inappropriate to directly apply the energy consumption data from this concentration process. The current data can not support the results presented in the article.", "reviewer_4": "1. Regarding the Capex of the CO2 electrolyzer, I still maintain that the authors are making a major mistake that results in significantly underestimating the value of the electrolyzer Capex. I see that they are using the same input numbers from a previous paper (Energy & Fuels, 2023, 37, 23, 17997–18008, Supplementary Information Page 12, Fig. R3), but when I looked deeper into this, it seems that the assumption of a reference current density of 0.17 A/cm2 in the Energy & Fuels paper is incorrect. It’s also worth pointing out that their Eqn. (6) is missing a term for the total cell area, another sloppy mistake.\nFrom page 4 of the cited 2019 NREL Report (“Manufacturing Cost Analysis for Proton Exchange Membrane Water Electrolyzers”, an excellent resource), one can clearly see the reference voltage of 1.7 V, and the current density should be 1.7 A/cm2, not 0.17 A/cm2. A current density of 1.7 A/cm2 is much more typical of a PEM water electrolyzer, which is the basis of the NREL study. From all previous electrolyzer TEA studies by others or that I have done myself, the electrolyzer Capex is based on normalization of the total electrolyzer system Capex by the rated power (i.e. Area x reference operating current density x reference voltage), not an arbitrarily low current density like 0.17 A/cm2. Why would a reference current density of 0.17 A/cm2 be chosen when the PEM electrolyzer typically operates 1.5-2.0 A/cm2? Thus, it seems evident that there is a mistake in the Energy & Fuels paper (and hence, the current manuscript) that leads to the CO2 electrolyzer Capex (expressed in $/m^2) being 1 order of magnitude lower than it should be.\nIn the pdf version of my comments, I've included the reference table from the 2019 NREL report cited in the Energy & Fuels paper that clearly shows a reference current density of 1.7 A/cm2, not 0.17 A/cm2.\n2. Regarding CO poisoning: my original point in my previous comment was that the CO2 electrolyzer is operating at such high voltages (4-5 V) that it is not surprising that there is no CO oxidation signal seen during CO stripping voltammetry measurement recorded at the end of the durability test. Under these conditions, the anode is likely polarized to such high values that any CO that crosses over from the cathode is easily oxidized and therefore doesn’t show up in the subsequent CO stripping voltammetry experiments. It may be true that CO poisoning isn’t occurring at these conditions, but that’s only because a very high cell voltage (and thus lower conversion efficiency) was used. The lack of a CO stripping peak after the stability test doesn’t prove that CO is not crossing over, and it doesn’t mean that CO poisoning is not an issue. I would guess that the use of a more CO tolerant anode catalyst like PtRu would lead to significantly reduced cell voltage. Thus, I still disagree with the authors’ statement that there is no CO crossover in their cell in this line: “ These results demonstrate that the CO generated at the cathode did not crossover to the anode to cause CO poisoning under the tested flow cell conditions.” At the least, I feel like this statement is very misleading and should be changed.", "reviewer_5": "I co-reviewed this manuscript with one of the reviewers who provided the listed reports. This is part of the Nature Communications initiative to facilitate training in peer review and to provide appropriate recognition for Early Career Researchers who co-review manuscripts."}, "author_response": "We sincerely appreciate the reviewer’s support. Following your suggestion, we have revised the cost associated with the concentration process based on both the reference1 methods and our experimental data. Regarding the CAPEX of the electrodialysis process, since the cost is highly related to the membrane area, we adjusted the membrane area according to our HCOOH concentration, as well as the electrodialysis current density and current efficiency reported in the reference2 to concentrate 0.1 M HCOOH. For the electrodialysis stack voltage, which depends strongly on the ionic conductivity of the electrolyte, we used a 0.5 M K2SO4-based electrolyte feed. As the input and output electrolyte conductivities in our system were similar to those in the reference1, we corrected only the OPEX by adjusting the energy consumption according to the current efficiency."}, {"round": 4, "reviewer_comments": {"reviewer_3": "The economic assessment of formic acid concentration in the current TEA is unreasonable and lacks detailed, solid information to support the accuracy of the adopted methodology.\nThe systems documented in the literature typically involve the separation of HCOONa solutions. In contrast, this study deals with a mixture of K2SO4 and HCOOH. These are entirely different reaction systems. Consequently, the energy consumption values reported in the literature cannot be directly applied to this specific mixture.\nIt is currently unclear to the reader how the electrodialysis method is implemented to concentrate the K2SO4/HCOOH mixture, particularly given the complication of anion crossover. Therefore, the reviewer finds the current calculation method to be unfounded.\nThe approach lacks support from both relevant literature and any provided experimental or simulation data. As a result, it is my opinion that the current methodology cannot reliably support the presented overall economic calculations.", "reviewer_4": "It appears that the updated TEA is now correct in terms of how it handles the CO2 electrolyzer Capex. This was my main remaining issue with this manuscript, which I believe is now suitable for publication.", "reviewer_5": "I co-reviewed this manuscript with one of the reviewers who provided the listed reports. This is part of the Nature Communications initiative to facilitate training in peer review and to provide appropriate recognition for Early Career Researchers who co-review manuscripts."}, "author_response": ""}]

Decision letter