Results and Discussion
Figure Figure11 shows DEMS experiments of CO2 reduction on a roughened polycrystalline gold electrode with a roughness factor of 20.3 in an electrolyte of 0.5 M NaClO4 + 1 mM HClO4. The top panels (A-C) show in black the measured CV when the electrolyte is purged with an Ar/CO2 mixture featuring CO2 partial pressures of 0.1 bar (A), 0.3 bar (B), and 0.5 bar (C), respectively. The middle panels below (D-F) show the ionic current for mass 2 as a function of the applied potential. The lower panels (G-I) show the ionic current for mass 28, which is proportional to the electrochemically evolved amount of CO.14 The mass spectroscopic traces were measured in parallel to the electrochemical trace.
The red curves in Figure Figure11A-C is the sum of the partial Faradaic currents of hydrogen and CO formation. Both quantities were determined from the ionic currents for masses 2 (shown in Figure Figure11D-F) and 28 (shown in Figure Figure11G-I) via eqs 1a and 1b, respectively. The Faradaic current expected from the sum of the partial Faradaic current of H2 and CO formation shows good overlap with the overall Faradaic current measured by the potentiostat below −0.5 V vs Ag|AgCl. That is, for all Ar/CO2 mixtures, the formed amounts of H2 and CO determined mass spectroscopically account for the charge passed in the potential region of CO2 reduction. This suggests that no other products than H2 and CO are formed, which is also supported by the absence of any signal in the ionic current for masses 16, 27, and 30, indicative of methane, ethylene, and acetaldehyde, respectively. Therefore, our results obtained in mildly acidic electrolytes do not differ from CO2 reduction at gold at higher pH.4,5
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As the reaction is conducted under steady flow, a constant current due to mass transport limited proton reduction is expected in Figure Figure11.6,7,12,15 This behavior is observed both in the Faradaic current (Figure Figure22A) and the ionic current for mass 2 (Figure Figure22B) when an argon-purged electrolyte of 0.5 M NaClO4 containing 1 mM HClO4 is used. The reduction process occurring when the potential is scanned below −0.6 V vs Ag|AgCl can be assigned exclusively to the evolution of hydrogen. This is not only indicated by the ionic current for mass 2 that mirrors the behavior of the Faradaic current but also dictated by logic as this is the only possible reduction reaction in the absence of other electroactive species in the electrolyte. Once a potential of −0.95 V vs Ag|AgCl is passed in the negative-going direction, hydrogen evolution enters a steady value, indicative of a mass transport limited current.
In Figure Figure11, a signal evolves in the ionic current for mass 2 as a potential of −0.6 V vs Ag|AgCl is passed. Hence, the presence of CO2 does not affect the onset potential of proton reduction. However, diffusion limitation as in Figure Figure22 is not achieved. This does not become evident from the CVs shown in Figure Figure11A-C, but from the ionic current for mass 2, which has a peak at −0.9 V vs Ag|AgCl and goes through a minimum at −1.28 V vs Ag|AgCl, the current of which decreases as the CO2 partial pressure increases. In parallel, the CO formation rate increases (i.e., ionic current for mass 28) with increasing CO2 partial pressure. In addition, the peak in the formation rate of hydrogen at −0.9 V vs Ag|AgCl coincides with the onset of CO formation at −0.85 V vs Ag|AgCl, suggesting that CO2 reduction suppresses hydrogen evolution. The same behavior is observed for electrolytes with lower proton concentration, as shown in Figures S1-S3 of the Supporting Information. Although neither CO nor H2 evolution enter diffusion limitation, the CVs in Figure Figure11 appear to feature a limiting current. This is due to the increasing partial current as a result of CO formation, which compensates the decrease due to decreasing H2 evolution.
As shown in reactions 2a and 2b, the reduction of CO2 to CO leaves an oxygen atom in oxidation state −II, which will react with protons or water to form water or two OH- ions.
In reaction 2a CO2 reacts with water instead of protons because we expect a local pH of 7 or higher at the electrode surface. That is, at the onset of CO2 reduction, the reduction of protons has nearly reached diffusion limitation, which indicates a local pH of close to 7. Because of the OH- formed in reaction 2a and because of water reduction, the local pH is bound to rise beyond 7 once the capacity of the CO2/HCO3- buffer is exhausted. Only in the very early stages of CO2 reduction, proton discharge is not yet fully diffusion limited, which leaves some protons for reaction 2b to proceed. Since OH- formed via reaction 2a diffuses away from the electrode surfaces, it can intercept protons before they reach the electrode (reaction 3), which are then no longer available to support hydrogen evolution.
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A similar mechanism was previously invoked to explain lack of hydrogen evolution in mildly acidic electrolytes parallel to oxygen reduction.16,17 The combination of reactions 2a and 3 is equivalent to reaction 2b, but with the important distinction that protons do not directly react with CO2 but rather with the OH- generated. Such a mechanism is in agreement with the experimental observation that CO2 reduction is pH-independent,18 that is, the relevant hydrogen donor for CO2 is water, not protons.19,20 The mechanistic interpretation for this observation is that CO2 is activated by electron transfer, decoupled from proton transfer, leading to a negatively charged (or polarized) CO2 intermediate, bound to the gold surface.21
In principle, the observations made in Figure Figure11 could also be interpreted as the competition of reduced CO2 species22 and protons for the same adsorption sites at the gold electrode.23 Following the argumentation of Chaplin and Wragg,23 adsorption of CO2 species, which becomes increasingly favorable as overpotential or CO2 partial pressure, blocks the adsorption of protons and therefore their reduction to H2. However, if proton reduction was a surface-limited process, the rate of proton reduction should decrease with decreasing roughness of the gold electrode. The opposite is observed in Figure S4, where the experiment of Figure Figure11 is repeated at a smooth gold electrode. With decreasing roughness factor, proton reduction increases, thus ruling out a mechanism based on competitive adsorption. As the CO formation rate is limited by the reaction kinetics, it scales with the true surface area and decreases with the electrode roughness. However, the availability of protons is limited by mass transport, which scales with the geometric surface area of the electrode. Although the flux of protons remains constant, less are used during CO formation, which leaves more protons for hydrogen evolution.
The competition for protons has a positive effect on the Faradaic efficiency for CO formation, which can reach 100% even in mildly acidic electrolytes. This is shown in Figure Figure33, which compares the Faradaic efficiency for CO formation as a function of potential for CO2 reduction from electrolytes with four different proton concentrations and three different CO2 partial pressures. The data from which the Faradaic efficiency in Figure Figure33 was calculated are shown in Figure Figure11 and Figures S1-S3 (Supporting Information). For any given proton concentration, the Faradaic efficiency increases with increasing CO2 partial pressure (panels A-C). This is due to both an increase of the CO formation rate and the increased suppression of hydrogen evolution. Of course, a higher proton concentration increases the rate of hydrogen evolution, and thereby lowers the Faradaic efficiency of CO formation. While the Faradaic efficiency of CO formation at the smooth electrode (c.f. Figure S5) shows the same behavior as that in Figure Figure33, the absolute values are consistently lower than those at the roughened gold electrode.
In Figure Figure33, the potential dependence of the Faradaic efficiency has a bell shape, which reaches its maximum in the potential range between −1.2 and −1.3 V vs Ag|AgCl, from which it decreases as the potential is made more negative.
This behavior can be understood from Figure Figure11D-F (as well as Figures S1-S3 in the Supporting Information), where the amounts of evolved hydrogen start to increase again as the potential is scanned below −1.28 V vs Ag|AgCl. In this potential region hydrogen evolution due to water reduction begins to take place,6,7 which is the dominant reason for the decreasing Faradaic efficiency. However, it also decreases because the CO formation rate either drops (Figure Figure11G,H) or its potential-dependent increase begins to flatten out (Figure Figure11I).
Closer inspection of Figures S1-S3 in the Supporting Information show that the formation rate of CO, that is, the partial Faradaic current due to CO formation, exceeds the decrease in the rate of proton discharge. That is, fewer protons are consumed during CO2 reduction than the CO formation rate suggests. In the potential range prior to water reduction, we can quantify the flux FCO2(H+) of protons that are consumed during CO2 reduction from eq 4
where If(2) is the partial Faradaic current of hydrogen evolution determined from the ionic current for mass 2, If,diff(2) is the diffusion-limited current of proton reduction when no CO2 reduction takes place, F is the Faraday constant, and z is the number of transferred electrons. The difference (If,diff(2) – If(2)) enters eq 4 because the current of protons diffusing to the surface (i.e., If,diff(2)) minus the protons reacting to hydrogen (i.e., If(2)) represent the current of protons that participate in a different reaction (i.e., CO2 reduction). In a similar way we can calculate from eq 5 the flux of CO formed during CO2 reduction (FCO2(CO))
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where If(28)represents the partial Faradaic current of CO formation. Furthermore, we can determine from the ionic current for mass 44 (c.f. Figure S6) the flux of CO2, F(CO2), that is consumed parallel to CO formation. Figure Figure44 shows FCO2(H+) divided by 2, FCO2 (CO), and F(CO2) as a function of the applied potential.
Since the formation of 1 molecule of CO consumes 1 molecule of CO2 and 2 protons, the black, red, and blue curves in Figure Figure44 should overlap. At low overpotentials this is indeed the case in Figures Figures44A-C, where because of the low reaction rate, the CO formation rate is low. However, at higher overpotentials in Figure Figure44B,C, where the formation rate of CO is higher because of the higher CO2 pressure, only a fraction of the consumed CO2 is actually reduced to CO. On the other hand, the number of protons participating in the reduction of CO2 is not large enough to support the observed CO formation rate. We can calculate this proton deficit, DH+, from eq 6:
and the surplus of CO2 consumption SCO2 from eq 7:
With increasing CO formation rate, both DH+ and SCO2 (lower panels of Figure Figure44) increase and overlap with good agreement. The same behavior is observed in Figures S7-S9 of the Supporting Information, which present the same data as Figure Figure44 for electrolytes with a proton concentration of 0.63, 0.4, and 0.25 mM, respectively. The fact that DH+ and SCO2 overlap quite well suggests that the proton deficit is compensated by CO2 forming bicarbonate near the electrode surface. That is, OH- formed during CO2 reduction via reaction 2a is not only neutralized by protons as in reaction 3 but also reacted with CO2 to form bicarbonate according to reaction 8.
The good overlap between that DH+ and SCO2 also means that we can completely account for the mass balance of CO2 consumption. That is, CO2 is consumed as a result of CO formation and the reaction with OH- formed during CO2 reduction as well as water reduction.
We show in the Supporting Information that SCO2 equals the rate of reaction 8 (i.e., rate of bicarbonate formation), whereas the consumption of protons during CO formation (i.e., FCO2(H+)) equals twice the rate of CO formation (rate of reaction 2a) minus the rate of bicarbonate formation (rate of reaction 8). This implies that complete suppression of proton reduction requires a CO formation rate that is equal to or exceeds the mass transport rate of protons to the electrode surface. Only then is the formation rate of OH-, which forms according to reaction 2a along with CO, sufficient to intercept all protons diffusing toward the electrode surface. We believe that this is a key guiding principle for designing an efficient electrolyzer for electrochemical CO2 reduction in acid media: it must accommodate such high CO2 reduction rates that the OH- formed as a byproduct can neutralize all protons that would otherwise participate in hydrogen evolution.
With reaction 8 we can also understand why we observe in Figure Figure11 that the CO formation rate flattens out or decreases as the potential decreases below −1.4 V vs Ag|AgCl. At this potential water reduction leads to the additional formation of OH-, resulting in the consumption of CO2. As shown in Figures S10-S12, this leads even under mass transport control to a significant drop of the CO2 partial pressure at the electrode surface. The potential-dependent increase of the rate constant of CO2 reduction cannot compensate for the effect of the decreasing local CO2 concentration. As a result, the current due to CO2 reduction drops. This is not limited to the potential region of water reduction. Also, in the potential region of proton reduction the CO formation rate increases slightly in Figures S1-S3 with increasing proton concentration. This is remarkable as protons are not involved in the rate-determining step, which means that their concentration should not affect the rate of CO2 reduction.22,18−21 However, a higher proton concentration means that a larger share of OH- formed during CO2 reduction is neutralized via reaction 3. Therefore, a higher proton concentration maintains a higher local CO2 concentration at the electrode surface and supports indirectly a higher CO formation rate.
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