Introduction

In order to alleviate greenhouse effects, countries around the world have formulated different carbon neutralization plans1. Electrochemical reduction of CO2 (CO2ER) can convert CO2 into valuable fuels and chemicals2,3,4,5,6,7, which is an emerging approach to utilize CO2 as a resource. CO2ER to either CO or formate through the two-electron transfer reaction is promising for industrialization due to its high selectivity and high current density8,9. An in-depth understanding of the two-electron transfer reaction mechanisms is, therefore, essential to designing high-efficiency catalysts and advancing the progress of CO2ER towards commercialization.

Recently, the two-electron transfer of CO2ER to formate and CO has been widely accepted to proceed through the pathways with intermediate species of *OOCH and *COOH, respectively10,11, as inner-sphere electron-transfer reactions12,13. However, there is still a debate in terms of the RLS14,15,16. Specifically, the RLS has been considered to be the adsorption of CO2 along with an electron transfer over the catalysts17,18,19,20,21, the transfer of protons to *CO2− to form *COOH or *OOCH16,22,23, the proton-coupled electron transfer to form *COOH or *OOCH24, or the desorption of *CO25,26 according to the results of Tafel analysis16,17,22, reaction order analysis20,21, and theoretical calculations25,26. These methods adopted previously may fail to recognize the correct RLS due to experimental errors or insufficient precision. For the Tafel analysis, it is hard to tell the RLS in CO2ER filed due to mass transport limitations16 and the value assumed for the transfer coefficient (α)27. Because the theoretical Tafel slope might be incorrect due to the improper assumption of 0.5 transfer coefficient which might lead to a wrong guidance for the RLS27. The reaction order analysis used to determine whether the RLS includes a proton transfer step is often performed by changing the pH of the electrolyte28,29. It is an incomplete approach because the source of protons may come from water molecules instead of H+[ 30, where the activity of CO2ER would not change with pH. On the other hand, the RLS without proton transfer may still change with pH if the proton transfer step occurs before the RLS. Although theoretical calculations have widely been adopted to find the RLS, the reliability of the results greatly depends on the choice of models and experimental methods31,32. Therefore, it is of great significance to also develop a reliable experimental strategy to reveal the RLS for CO2ER.

In this work, theoretical derivations and experimental results are combined and discussed in detail to provide sufficient evidence for the determination of the RLS during the two-electron transfer CO2ER. The rate expressions with different RLSs were first derived through the breakdown of the Butler-Volmer equation. By analyzing these rate expressions, an effective method for discovering the RLS was proposed, which combines pH dependency and kinetic isotope effect (KIE) experiments. Since H+ ions and H2O molecules are both proton donors, the reaction orders of H+ and H2O can be obtained by changing the pH of the electrolyte and the adoption of D2O in the electrolyte. Corresponding results could help clarify whether the adsorption of CO2 with its concomitant electron transfer (ET) step or the other possible steps, such as the proton transfer (PT), proton-coupled electron transfer (PCET), or desorption (D) of product, is the RLS. Subsequently, Au, Ag, Sn, and In were used as model catalysts to reveal the RLS of the two-electron transfer CO2ER. For all the electrocatalysts, the current densities of CO (jCO) and formate (jHCOO−) are independent of both pH and deuteration of the electrolyte, which indicates the CO2 adsorption step to be the RLS.

Results

The rate expressions with different reaction steps as the RLS

In order to discover the RLS, the Butler-Volmer equation was employed to describe the kinetic rate expression of two-electron transfer CO2ER. It describes how the electrical current passing through an electrode depends on the voltage difference between the electrode and the bulk electrolyte for simple unimolecular redox reactions, when both a cathodic and an anodic reaction proceeding on the same electrode are controlled by surface reactions rather than the mass transfer of electrolyte33. For electroreduction reactions (Eq. 1, where Ox and Red represent oxidant and reductant, respectively), the Butler-Volmer equation is shown as Eq. 234.

$${{{{{\rm{O}}}}}}_{{{{{\rm{x}}}}}}+{{ne}}^{-}\to {{{{{\rm{R}}}}}}_{{{{{\rm{ed}}}}}}$$
(1)
$$j=nF{k}_{b}^{0}a[{{{{{{\rm{R}}}}}}}_{{{{{{\rm{ed}}}}}}}]\exp [(1 - \alpha )f\eta ] - {{{{{\rm{n}}}}}}F{k}_{f}^{0}a[{{{{{{\rm{O}}}}}}}_{{{{{{\rm{x}}}}}}}]\exp [ - \alpha f\eta ]$$
(2)

In Eq. 2, j is the current density; η is the overpotential for the cathodic reaction; kf0 is the standard forward rate constant; kb0 is the standard backward rate constant; F is the Faraday constant; f = F/RT, where R is the ideal gas constant and T is absolute temperature; α is the transfer coefficient assumed to be equal to 0.5; n is the number of transferred electrons; a[Red] and a[Ox] are the concentrations of reductant and oxidant.

When the overpotential is sufficiently high, i.e., exp[–(1–α)f\(\eta\)] << exp(–αf\(\eta\)), the backward reaction can be ignored35. Even the high-performance CO2 electrolysis catalysts have sufficient overpotentials to meet this condition36. Therefore, Eq. 2 can be simplified to Eq. 3. At equilibrium conditions (j = 0), Eq. 2 can be simplified to Eq. 4.

$$j= - nF{k}_{f}^{0}a[{{{{{{\rm{O}}}}}}}_{{{{{{\rm{x}}}}}}}]\exp ( - \alpha f\eta )$$
(3)
$$a[{{{{{{\rm{R}}}}}}}_{{{{{{\rm{ed}}}}}}}]/a[{{{{{{\rm{O}}}}}}}_{{{{{{\rm{x}}}}}}}]={K}^{\theta }\exp ( - f\eta )$$
(4)

By combining Eq. 4, the a[Ox] in Eq. 3 can be represented by the concentration of reactants and Kθexp(–f\(\eta\)) in the previous step (see the supplementary information for more details). Subsequently, the rate expression of the two-electron transfer CO2ER with a specific reaction step as the RLS can be derived (Tables 1 and 2, different labels are assigned to the corresponding RLSs according to the reaction processes). Whether the RLS is controlled by ET, PT, PCET, or D is also shown in the Tables. One thing should also be kept in mind is that all these expressions are based on assumptions of what might happen in the mechanism, which may not cover all possible kinetic cases at current cognitive levels.

Table 1 Reaction kinetic parameters for different possible RLSs during the two-electron transfer CO2ER to COa.
Full size table
Table 2 Reaction kinetic parameters for different possible RLSs during the two-electron transfer CO2ER to formatea.
Full size table

According to the rate expression, the reaction order of different reactants can be obtained. For example, when the adsorption of CO2 with the ET (step A1 in Table 1, Eq. 5) is the RLS for CO2ER to CO,

$${{{{{{\rm{CO}}}}}}}_{2}+\ast +{e}^{-}\to \ast {{{{{\rm{CO}}}}}}_{2}^{-}$$
(5)

the rate expression (Eq. 6) is

$${j}_{{{{{\rm{co}}}}}}=2F{k}_{A1}^{0}a[{{{{{{\rm{CO}}}}}}}_{2}]{\theta }^{\ast }\exp ( - \alpha f\eta ),$$
(6)

where the corresponding reaction order of H+ and H2O molecules should be 0.

Thus, the RLS of two-electron transfer CO2ER could be determined via the analysis of the reaction order of the reactants19,28,30,37. Whether the reaction is controlled by the concentration of H+ can be reflected by its pH dependency. However, whether protons are involved in the RLS cannot be simply determined by the pH dependency of the reaction, since H2O could be the proton source. Therefore, KIE experiments can be conducted to reveal if the H2O molecules are involved in the reaction as a proton source. It is noteworthy that possible RLS with the rate expression involving neither H+ nor H2O may still be controlled by them because protons may take part in the reaction processes before the RLS (see the supplementary information for more details). Therefore, in-depth reaction rate analysis is essential for the determination of the RLS.

The fabrication of model catalysts

To elucidate the RLS of two-electron transfer CO2ER, Au, Ag, Sn, and In were chosen as model catalysts. Au and Ag have been proved to exhibit good performance for the production of CO, and In and Sn are promising catalysts with high selectivity for formate38,39. These catalysts were deposited on Si(100) wafers by magnetron sputtering. To enhance the adhesion between the catalysts and the Si wafers, Ti films with a thickness of approximately 15 nm were first deposited on the Si wafers40. Catalysts with relatively high conductivity (i.e., Au and Ag) were directly deposited on the Ti films. The thicknesses of the catalyst films were controlled to be 200 nm. Less conductive In and Sn catalyst films with a thickness of about 400 nm were deposited after the adhesion of 30 nm Au layers onto the Ti films. This strategy improves the conductivity of the substrate and preventes delamination of the films under cathodic potentials. Because of the high surface tension of Sn and In, thicker films need to be deposited to cover the substrates completely. According to scanning electron microscopy (SEM) images (Supplementary Fig. 1), the Au, Ag, In, and Sn catalyst films are evenly distributed over the substrates. X-ray diffraction (XRD) patterns (Supplementary Fig. 2) show that these films have polycrystalline structures. No signal of the substrate materials was found in the survey X-ray photoelectron spectroscopy (XPS) spectra of the samples, indicating that no substrate would be exposed to the electrolyte (Supplementary Fig. 3). Furthermore, the analyses of the surface valence states of the Au and Ag films show that they are primarily in the metallic state with only slight surface oxidation for Ag. However, the surfaces of the Sn and In films were oxidized (Supplementary Fig. 4). These four model catalysts were then used to study the RLS of two-electron transfer CO2ER. Here, CO and H2 were the main products over Ag and Au. CO, formate, and H2 were the main products over Sn and In. The total Faradaic efficiency is basically equal to 100% (Supplementary Fig. 5).

Experimental determination of the RLS

In order to explore the impact of the H+ concentration on the two-electron transfer CO2ER, the changing trends of jCO and jHCOO− with the variation of electrolyte pH were analyzed. Experiments were carried out in a flowing H-cell with five types of CO2 saturated electrolytes (0.3 M KHCO3, pH 7.0; 0.1 M K3PO4, pH 6.6; 0.1 M KH2PO4, pH 4.3; 0.1 M KH2PO4 + 0.1 M H3PO4, pH 2.9; 0.1 M H3PO4, pH 1.6). Figure 1a–d show the jCO of the four catalysts as a function of applied potential. The electrocatalytic activity for CO generation is barely affected by the pH of the electrolytes (Supplementary Fig. 6a). Since the jCO is consistent under various pH from 2.9 to 7.0, the reaction rate expression of CO2ER to CO should not include H+. Therefore, only the following reaction steps (Eqs. 7–9), with reaction order for H+ to be 0 in the rate expressions, could possibly be the RLS.

$${{{{{\rm{A}}}}}}1\,{{{{{\rm{or}}}}}}\,{{{{{\rm{a}}}}}}1:{{{{{{\rm{CO}}}}}}}_{2}+\ast +{{{{{{\rm{e}}}}}}}^{-}\to \ast {{{{{{{\rm{CO}}}}}}}_{2}}^{-}$$
(7)
$${{{{{\rm{a}}}}}}2:\ast {{{{{{{\rm{CO}}}}}}}_{2}}^{-}+{{{{{{\rm{H}}}}}}}_{2}{{{{{\rm{O}}}}}}\to \ast {{{{{\rm{COOH}}}}}}+{{{{{{\rm{OH}}}}}}}^{-}$$
(8)
$${{{{{\rm{b}}}}}}1\,{{{{{\rm{or}}}}}}\,{{{{{\rm{c}}}}}}1:{{{{{{\rm{CO}}}}}}}_{2}+\ast +{e}^{-}+{{{{{{\rm{H}}}}}}}_{2}{{{{{\rm{O}}}}}}\to \ast {{{{{\rm{COOH}}}}}}+{{{{{{\rm{OH}}}}}}}^{-}$$
(9)
Fig. 1: The pH dependency for CO2ER toward CO and formate.
figure 1

The jCO for Au (a), Ag (b), Sn (c), and In (d) catalysts in electrolytes with different pH. The jHCOO− for Sn (e) and In (f) catalysts in electrolytes with different pH. Error bars are means ± standard deviation (n = 3 replicates).

Full size image

Similar to the case of CO, the jHCOO− of both Sn and In did not show significant change with the electrolyte pH (Fig. 1e, f and Supplementary Fig. 6b). Thus, the reaction order of H+ for CO2ER to formate should also be 0. According to Table 2, the possible RLS meeting this requirement are listed as follows (Eqs. 10–12).

$${{{{{\rm{D}}}}}}1\,{{{{{\rm{or}}}}}}\,{{{{{\rm{d}}}}}}1:{{{{{{\rm{CO}}}}}}}_{2}+\ast +{{{{{{\rm{e}}}}}}}^{-}\to \ast {{{{{{{\rm{CO}}}}}}}_{2}}^{-}$$
(10)
$${{{{{\rm{d}}}}}}2:\ast {{{{{{{\rm{CO}}}}}}}_{2}}^{-}+{{{{{{\rm{H}}}}}}}_{2}{{{{{\rm{O}}}}}}\to \ast {{{{{\rm{OOCH}}}}}}+{{{{{{\rm{OH}}}}}}}^{ - }$$
(11)
$${{{{{\rm{e}}}}}}1:{{{{{{\rm{CO}}}}}}}_{2}+\ast +{{{{{{\rm{e}}}}}}}^{ - }+{{{{{{\rm{H}}}}}}}_{2}{{{{{\rm{O}}}}}}\to \ast {{{{{\rm{OOCH}}}}}}+{{{{{{\rm{OH}}}}}}}^{ - }$$
(12)

On the contrary, the activity of the hydrogen evolution reaction (HER) is enhanced as the pH decreases (Fig. 2 and Supplementary Fig. 7). This phenomenon indicates that the RLS of HER depends on the concentration of H+, which is consistent with the results from the literatures30,41. This result further supports the feasibility of revealing the RLS of two-electron transfer CO2ER by pH dependency.

Fig. 2: The pH dependency for HER.
figure 2

The jH2 for Au (a), Ag (b), Sn (c), and In (d) catalysts in electrolytes with different pH. Error bars are means ± standard deviation (n = 3 replicates).

Full size image

Considering the inevitable deviation of the bulk pH and local pH caused by the constant consumption of protons during CO2ER and HER, the local pH was investigated. Supplementary Fig. 9 shows the local pH as a function of the operating potential for In (see the note of Supplementary Fig. 9 for the reason why choosing In as the representative catalyst). The local pH elevates slowly with the increase of the reaction potential in different electrolytes. However, the order of the local pH is the same as that of the bulk pH. Therefore, it is reasonable to use bulk pH in the current work. Also, regarding the inevitable change of ion concentration when changing pH, comparative experiments in 0.1 and 0.3 M KH2PO4 show the jCO and jHCOO− are basically the same (Supplementary Fig. 10), indicating the change in the concentrations of potassium and phosphate has little effect on the CO2ER activity in those experiments. So, the ion concentration caused by changing the pH will not affect the conclusion.

The pH dependency experiments eliminated many potential RLSs. To further determine the RLS, KIE experiments were considered to analyze whether water molecules are involved in the rate expression. Since the KIE of CO2ER in homogeneous catalysts was tested to be 6.9242 and 8.243 for CO and formate, respectively, which means if proton was involved in the RLS for CO2ER to CO or formate, the KIE should be >1. Secondly, the KIE was used to exclude the step of H2O providing proton, where the KIE of H2O dissociation are around 3.4–7.644, so the involvement of protons in the RLS of CO2ER should show the KIE>1. Thirdly, KIE experiments have been known to lead to false-negative conclusions, but only in very specific and rare instancesnormally involving more environment-sensitive molecular catalysts45. To the best of our knowledge, there have been no calculation on transition-metal catalysts showing D2O could distort CO2 electrolysis results for two-electron products compared to using H2O. Therefore, it seems reasonable to use KIE to explore whether water molecules are involved in the rate expression.

In the KIE experiments, both jCO and jHCOO− did not change with the use of D2O instead of H2O in the 0.1 M KH2PO4 electrolyte for Au, Ag, Sn, and In catalysts (Fig. 3). To eliminate the possibility that KH2PO4 provides protons that equilibrate with D2O to generate a small amount of H2O, K2CO3 solutions in H2O and D2O were also chosen as electrolytes. The results of activity tests with the In catalyst show that jCO and jHCOO− are almost the same in both electrolytes (Supplementary Fig. 11a,b). These phenomena demonstrate that the reaction order of H2O for CO2ER to CO or formate should be 0. Combining this knowledge with the results of pH dependency studies, the RLS of both the two-electron transfer CO2ER to CO and formate is deduced to be the adsorption of CO2 with one electron transferred simultaneously, as shown in Eqs. 7 and 10, respectively. KIE experiments for HER were also conducted with the Au, Ag, Sn, and In catalysts. When D2O was used in the solvent, the jH2 drops significantly (Fig. 4 and Supplementary Fig. 11c), indicating water molecules are a part of the RLS or take part in the reaction process before the RLS, which is also consistent with previous reports46. It should be noticed that the pH dependency and KIE experiments were not conducted in alkaline electrolytes since the solution would become neutral or acidic after the dissolution of CO247, but we think the conclusion might also be applicable to alkaline solutions (see Supplementary Figs. 12 and 13 for detailed explanation).

Fig. 3: Kinetic isotope effect for CO2ER toward CO and formate.
figure 3

The jCO for Au (a), Ag (b), Sn (c), and In (d) in 0.1 M KH2PO4 electrolytes with D2O and H2O. The jHCOO− for Sn (e) and In (f) in 0.1 M KH2PO4 electrolytes with D2O and H2O. Error bars are means ± standard deviation (n = 3 replicates).

Full size image
Fig. 4: Kinetic isotope effect for HER.
figure 4

The jH2 for Au (a), Ag (b), Sn (c) and In (d). The data were obtained from running CO2ER experiments at 0.1 M KH2PO4 electrolyte with D2O and H2O. Error bars are means ± standard deviation (n = 3 replicates).

Full size image

Based on the above analyses, the RLS of two-electron transfer CO2ER for Au, Ag, Sn, and In catalysts was found to be the adsorption of CO2 with the concomitant ET. Since the transfer of electrons is quite fast over metal catalysts28,48, the ultimate RLS is proposed to be most likely the adsorption of CO2.

Discussion

In summary, this work presents an effective approach to determine the RLS by employing a detailed analysis of the reaction rate expression together with pH dependency and KIE experiments. It was found that both jCO and jHCOO− are independent of the pH and deuteration of the electrolytes for Au, Ag, Sn, and In, which are representative catalysts for two-electron transfer CO2ER. The results reveal that the RLS of the two-electron CO2ER should be the adsorption of CO2. This finding suggests effective strategies to design highly active CO2ER catalyst for the production of CO and formate.

Methods

Electrode preparation

Au, Ag, In, and Sn thin films were deposited onto single-crystal Si wafers with the (100) orientation using an AJA ATC Orion-5 magnetron sputtering system. Before the deposition, the Si wafers were etched with Ar+ ions for 5 min with a power of 40 W to clean the silicon oxide on Si wafers. Then, 15 nm Ti films were deposited as binders between catalysts and Si wafers at the power of 130 W. Au and Ag catalyst films with the thickness of about 200 nm were deposited over Ti at 50 W. Due to the poor electrical conductivity of In and Sn catalysts, Au films (30 nm) were added over Ti before depositing 400 nm of the catalyst films at 20 W.

Electrode characterization

The crystal structures of the Au, Ag, In, and Sn thin films were analyzed with a Rigaku Smartlab X-ray diffractometer (XRD) using Cu Kα radiation (40 kV, 40 mA). The near-surface compositions of the thin films were measured with a Kratos Axis Ultra DLD X-ray photoelectron spectrometer (XPS). All spectra were acquired using monochromatized Al Kα radiation (15 kV, 15 mA). The kinetic energy scale of the measured spectra was calibrated by setting the C 1s binding energy to 284.8 eV. The surface structure of those thin films was recorded using an FEI XL30 Sirion scanning electron microscope (SEM) at the 5 kV acceleration voltage, Everhart-Thornley detector, secondary electrons mode.

Electrochemical characterization

All electrochemical activity measurements were conducted in a custom electrochemical cell machined from PEEK at room temperature and atmospheric pressure. The cell was sonicated in 20 wt% nitric acid and thoroughly rinsed with DI water before all experimentation. The working and counter electrodes were parallel and separated by a bipolar membrane (Fumasep FBM). The exposed geometric surface area of each electrode was 1 cm2. The electrolyte volumes in the cathodic and anodic chambers were 6 mL and 1 mL, respectively. The counter electrode was iridium dioxide (IrO2) purchased from Dioxide Materials. The working electrode potential was referenced against a miniature Ag/AgCl electrode (Innovative Instruments Inc.) that was calibrated against a homemade standard hydrogen electrode (SHE). 0.3 M potassium bicarbonate (KHCO3, Sigma Aldrich 99.7%, pH 7.0), 0.1 M potassium phosphate (K3PO4, Sigma Aldrich 99.99%, pH 6.6), 0.1 M monopotassium phosphate (KH2PO4, Sigma Aldrich 99.99%, pH 4.3), 0.1 M monopotassium phosphate adjusted the pH with 0.1 phosphoric acid (KH2PO3 with H3PO4, pH 2.9) and 0.1 M phosphoric acid (H3PO4, Sigma Aldrich 85% w/w, pH 1.6) solutions prepared using 18.2 MΩ·cm Milli-Q water were used as the cathodic electrolyte. 0.1 M KH2PO4 was used as the anodic electrolyte. Metallic impurities in the as-prepared electrolytes were removed before electrolysis by chelating them with Chelex 100 (Sigma Aldrich). The cathodic electrolyte was sparged with CO2 (99.999% Praxair Inc.) at a rate of 10 sccm for 30 min prior to the experiments. Then CO2 was pumped into the cathodic chamber by using a peristaltic pump (SHENCHEN LabN6) with the rate of 130 rev/min. Here, the experiments process in this pump speed cannot be the significantly diffusion-limited, since we have obtained relatively straight Tafel slopes over three orders of magnitude in current (Figs. 1–4). The pH values of electrolytes were measured in CO2 saturated solutions.

The produced CO and H2 were tested by gas chromatography (GC, Thermo scientific, TRACE 1300). Ar was used as the carrier gas. The GC was equipped with a packed Molsieve 5A column, a packed Hayesep Q column, and a Rt-Qbond column to separate the gaseous products. Thus, H2 and CO could be identified using a thermal conductivity detector and a flame ionization detector, respectively. The liquid-phase products were analyzed after the electrolysis using a high-performance liquid chromatography (HPLC, Agilent 1200 series). Liquid-phase products were separated by an Aminex HPX-87H column (Bio-Rad) that was maintained at 50 °C. The HPLC was equipped with a diode array detector (DAD) and a refractive index detector (RID). The response signals of the DAD and RID were calibrated by solutions with different concentrations.

Electrochemical characterizations were performed using a Biologic VSP-300 potentiostat. All electrochemical measurements were recorded versus the reference electrode and converted to the SHE scale. Current interrupt was used to determine the uncompensated resistance (Ru) of the electrochemical cell. The accurate potentials were corrected according to the Ru (see Supplementary Table 3 for detailed potential correction process).

The electrocatalytic activity of the thin films was assessed by conducting chronoamperometry with a step length of 10 min. Each thin film was tested at least three separate times to ensure statistical relevance of the observed trends. The Faradaic efficiency and partial current density calculation process can be found in Supplementary information.