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COMMUNICATION
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How copper catalyzes the electroreduction of carbon dioxide into hydrocarbon
fuels†
Andrew A. Peterson, Frank Abild-Pedersen, Felix Studt, Jan Rossmeisl and Jens K. Nørskov*
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Published on 26 August 2010 on http://pubs.rsc.org | doi:10.1039/C0EE00071J
Received 10th May 2010, Accepted 17th June 2010
DOI: 10.1039/c0ee00071j
Density functional theory calculations explain copper’s unique
ability to convert CO2 into hydrocarbons, which may open up
(photo-)electrochemical routes to fuels.
The storage of energy in chemical bonds, as fuels, is attractive for
a number of reasons,1 and a (photo-)electrochemical route to reduce
CO2 to hydrocarbon fuels would provide an ideal storage medium for
intermittent renewable energy sources, resulting in carbon-neutral
fuels. However, no material is known to catalyze the electroreduction
of CO2 to fuels both efficiently and selectively. In fact, only copper
and its alloys have been shown to be capable of producing significant
quantities of hydrocarbons from CO2, but they do so inefficiently
with a large overpotential requirement. In this communication, we
use a computational hydrogen electrode (CHE) model to show how
copper is able to catalyze this reaction, and we outline requirements
for more efficient catalysts to enable artificial photosynthesis.
Numerous researchers2–10 have studied the electrochemical reduction of CO2 at metal electrodes, and two excellent reviews have
appeared in the recent literature.11,12 Copper has been found to be
unique among the metals in its ability to produce a high quantity of
hydrocarbon fuels from the electroreduction of CO2.2 A typical
product distribution as a function of potential, as measured by Hori
et al.7 with a Cu electrode, is shown in Fig. 1. The hydrocarbons
methane (CH4) and ethylene (C2H4) are the dominant products at
sufficiently negative potentials. At less negative potentials, hydrogen
(H2), formic acid (HCOOH), and carbon monoxide (CO) are instead
dominant. It is interesting that methanol (CH3OH) is not among the
products reported since copper catalysts are commonly used to
Center for Atomic-scale Materials Design, Department of Physics,
Technical University of Denmark, DK-2800 Lyngby, Denmark
† Electronic supplementary information (ESI) available: Computational
details. See DOI: 10.1039/c0ee00071j
Fig. 1 Experimentally determined current and product distribution at
a copper electrode. Product distribution and total current produced as
a function of applied potential (versus reversible hydrogen electrode,
RHE) in the electrochemical reduction of CO2 at a copper electrode in
0.1 M KHCO3 (pH 6.8) at 18.5 C, as measured by Hori et al.7
produce methanol selectively from a mixture of CO2, CO, and H2 in
the methanol synthesis reaction.13
Although copper has been shown to be unique in producing
hydrocarbons from CO2, it is remarkably inefficient in doing so.
Thermodynamically, a potential of +0.17 V vs. RHE is all that is
required for the reaction:
Broader context
Hydrocarbon fuels provide unparalleled energy density and are the backbone of our energy infrastructure. However, today’s only
sources of hydrocarbons are fossil fuels and biomass. If an efficient electrochemical process could be developed to produce
hydrocarbons from CO2, this could allow processes (known as ‘‘artificial photosynthesis’’ or ‘‘solar fuels’’) to directly produce useful
fuels from CO2 and renewable energy sources, which are often intermittent and will require storage when deployed at large scales. It
was first shown more than 20 years ago that copper can act as an electrocatalyst in reducing CO2 into the hydrocarbons CH4 and
C2H4. Copper is unique in its ability to perform this catalysis, but it does so with a relatively high overpotential (1 V). The
mechanism by which copper performs this catalysis, as well as the reason for the overpotential are unknown. The current
communication outlines a plausible mechanism, based on quantum chemical simulations, that explains the experimental observations and may help define design principles for improved electrocatalysts.
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CO2 + 8(H+ + e) / CH4 + 2H2O
(1)
at 18.5 C. However, in experiments, potentials of about 0.8 V are
required for the onset of CH4 production from CO2, and 1.0 V is
required for a decent current (to CH4) of 2 mA cm2.7
The mechanisms by which copper produces this unique product
distribution and the reason for the overpotential have remained
elusive. This opens up three main questions to which we propose
answers in this work. (1) How does copper catalyse this unique
product spectrum, including hydrocarbons, from CO2? (2) Why is
such a large overpotential required? (3) Why does the pathway
produce methane, but not methanol, given that methanol is the
dominant product on copper catalysts in industrial processes? These
answers should shed light on the true barriers to making artificial
photosynthesis a commercial reality.
It is extremely demanding to provide a detailed theoretical
description of chemical processes taking place at an electrified solidwater interface.14–25 However, theoretical techniques have recently
been developed that provide an elegant method26 of modeling electrochemical reactions using what we will herein refer to as a computational hydrogen electrode (CHE) model; this model is coupled with
adsorption energies from density functional theory (DFT) calculations. Using these techniques, electrochemical reaction pathways can
be elucidated and the voltage requirements at which different
chemical pathways open can be estimated. We will show that by
using the simple CHE model as a starting point, the major trends in
CO2 reduction over Cu surfaces can be elucidated.
The CHE model was applied to the electroreduction of CO2 by
examining a network involving 41 different intermediate steps on the
Cu (211) surface. Full details, including a description of the CHE
model, are available in the Supplementary Information. From this
pathway, numerous routes to the major products CO, H2, HCOOH,
CH4 and C2H4 are possible and the lowest energy pathways were
found as a function of applied potential.
The least-negative potential at which the pathway to each product
becomes exergonic (downhill in free energy) is referred to as the
limiting potential, which serves as a first estimate of the onset
potential for each species. More exact predictions of onset potentials
will require calculations not just of free energies, but of barriers
between steps along the pathway. However, each electrochemical step
discussed in this communication involves the transfer of a proton
from solution to an adsorbed species on the surface. Barriers for
electrochemical proton transfers have been calculated for the reduction of O2 to OOH on Pt14 and for the reduction of OH to H2O on
Pt.27 In both cases, the proton-transfer reaction barriers were calculated to be small (0.15 eV to 0.25 eV) at the potential needed to make
the elementary step exergonic, and were found to diminish with
higher applied voltages. Similarly, as a first approximation we expect
that barriers for electrochemical proton transfers to adsorbed species
in this study will be small and easily surmountable at room temperature. However, the same assumption cannot be made for reactions
between two adsorbed species, which will be important in ethylene
formation (which involves the formation of a C–C bond) and in
a competing, non-electrochemical pathway to methanol, both of
which will be discussed later in this communication.
Fig. 2 shows free energy diagrams for the lowest-energy pathways
for the formation of H2, HCOOH, CO, and CH4 from CO2. In this
figure each zone (on the x axis) represents the transfer of one protonelectron pair. The free energy pathways at 0 V (vs RHE) are shown in
1312 | Energy Environ. Sci., 2010, 3, 1311–1315
black. Since each successive zone contains the chemical potential of
one extra proton-electron pair, the free energy (DGn) of each intermediate will change as a simple linear function of the applied
potential (U):
DGn(U) ¼ DGn(U ¼ 0) + neU
where n is the number of proton-electron pairs transferred relative to
CO2 and e is the elementary (positive) charge. By simulating more
negative voltages in the CHE model, the voltage at which different
chemical pathways become exergonic can be generated. These are
illustrated by the red pathways in Fig. 2.
Part (a) of Fig. 2 shows the first pathway to open in this system, in
which only H2 is produced; CO2 does not participate in the reaction.
The adsorption of the first proton from solution (State 1 to 33) is the
last step to become exergonic as a function of applied potential; this is
the potential-limiting step. The CHE model indicates that this is the
first pathway to open, consistent with the results in Hori et al.’s
experimental results reproduced in Fig. 1. This reaction pathway is
predicted by the CHE model to open on the Cu (211) surface at
around 0.03 V, but in experiments, significant current was not
observed until more negative potentials. (In Fig. 1, the first current is
reported at about 0.4 V, which goes almost entirely to produce H2.)
This can be understood with the CHE model: oxidation of the Cu
(211) step sites would block the production of H2. According to the
CHE model, O and OH would bind to the same stepped sites as H
and a potential of 0.3 V would be necessary to clear adsorbed OH
from these Cu step sites, enabling H2 evolution.
Fig. 2(b) shows the reaction pathway to formic acid (HCOOH)
predicted with the CHE model. Consistent with the experimental
results, this is the second pathway to open. In this pathway, CO2 and
a proton-electron pair adsorb as a carboxyl species; the addition of
a second proton-electron pair to this adsorbate results in the
production of HCOOH. The potential-limiting step is the formation
of carboxyl, requiring 0.41 V by the CHE model. This potential is
consistent with the onset in Fig. 1. A second pathway to formic acid
also exists (not shown in the figure), in which CO2 adsorbs as formate
(OCHO). This may be responsible for some of the formic acid
production at more negative potentials; the desorption of formate as
HCOOH is calculated to require 0.61 V to become exergonic.
Opening simultaneously with the HCOOH pathway in the
experimental results of Fig. 1 is that to CO, and this is also the next
pathway in the CHE model, opening at 0.41 V, which is illustrated
in Fig. 2(c). This route is also limited by the formation of carboxyl.
After this species adsorbs, the free-energy pathway is downhill to
remove water and produce adsorbed CO. This CO is weakly bound
at these conditions, leading to the production of CO both as a gas and
a surface species.
According to the CHE model, the key step in the formation of the
hydrocarbons CH4 and C2H4 is the hydrogenation of the adsorbed
CO to form adsorbed CHO at 0.74 V. From this point, the
hydrocarbon-forming reaction pathways open.
The lowest-energy pathway to methane is shown in Fig. 2(d). After
the formation of adsorbed CHO, a second proton-electron pair is
added to form adsorbed formaldehyde (H2CO), which binds weakly
to the Cu (211) step.28 The formaldehyde is protonated to form
adsorbed methoxy (OCH3), which stands up in an orientation with its
methyl group away from the surface. At this stage, a proton from
solution takes the methyl group off the methoxy, producing CH4.
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Fig. 2 Calculated free energy diagrams. Free energy diagrams for the lowest energy pathways to (a) H2, (b) HCOOH, (c) CO, and (d) CH4. In each
diagram, the black (higher) pathway represents the free energy at 0 V vs. RHE and the red (lower) pathway the free energy at the indicated potential. The
numbers labeled on the states correspond to the labels in Figure S1 of the Supplementary Information.
The remaining adsorbed oxygen is subsequently cleared as water.
After the production of adsorbed CHO, the formation of CHOH on
the surface is possible but less thermodynamically stable than the
route highlighted; this is a second potential route to CH4 that will not
affect the potential dependence, but will be of importance in future
kinetics studies.
The CHE model suggests several pathways to ethylene (C2H4), all
of which proceed through the CHO intermediate. Thus, protonation
of CO to form adsorbed CHO will be the potential-limiting step,
predicting these pathways to open at about 0.74 V. Ethylene
formation requires the creation of a C–C bond, which will occur in
a non-electrochemical step on the surface. Reaction free energies of
these C–C bond formation steps are summarized in Table 1; as
shown, all of the table entries are exergonic. Since these surface
reactions are not simple proton-transfer steps, the CHE model presented does not distinguish between them, and a future model that
incorporates reaction barriers will be required to determine which
pathway is dominant for C2H4 formation.
The experimental results show that H2 production decreases as the
production of CO and hydrocarbons is increased, including an
interesting dip in the total current around 0.8 V (see Fig. 1). We can
understand these effects as a result of changes in surface coverages.
As CO production becomes possible, adsorbed CO will start blocking
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Table 1 Calculated reaction free energies for C–C bond formation at
18.5 C. An asterisk (*) indicates that the species is adsorbed on the
surface
DGrxn
2 CH2O* / 2 O* + C2H4
CHO* + *OCH3 / 2 O* + C2H4
CH2O* + CHO* / *OCHCH2 + O*
2 CHO* / *OCHCHO*
0.73 eV
0.20 eV
1.04 eVa
1.12 eVa
a
The subsequent electrochemical proton addition steps to form C2H4
from *OCHCH2 and *OCHCHO* are exergonic at 0.74 V.
the surface sites, decreasing the number of active sites available. This
results in decreased production of H2 and a decreased overall Faradaic current. When the route to protonate the adsorbed CO becomes
exergonic, the CO clears and total current recovers.
We have shown that our analysis is in good agreement with CO2
electroreduction experiments.7 Other experimental observations
confirm the feasibility of the pathways presented here. Experiments
starting with HCOOH showed no detectable products,5,7 while
studies employing a CO feedstock have shown CH4 and C2H4
production.3,4,6 This validates the reactions in the early stages of the
mechanism, in that the hydrocarbon-generating reaction proceeds
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through CO, whereas the formate pathway is a dead end. (According
to the present model, adding an additional proton-electron pair to
form adsorbed OCH2O from OCHO would require 1.9 V vs. RHE
be applied, rendering this pathway unfeasible at the voltages reported.) Other studies5,6 have shown CH4 formation from a starting
material of CH2O, which is in agreement with the downstream
pathways shown in Fig. 2(d). Additionally, in experiments starting
with CO, H2 production was suppressed at a less negative potential
than when starting with CO2.7 This again is consistent with the CHE
model, which predicts a barrier for the conversion of CO2 to CO,
resulting from the adsorption of the COOH species. As evidenced in
Fig. 2(c), the adsorption of CO to the surface is not a function of the
applied potential, but around 0.41 V is required in order to get CO2
to produce CO.
It is worth considering why methanol (CH3OH) is formed in gasphase chemistry13 and methane (CH4) is formed in electrochemistry7,8
on copper surfaces. In the current study, we considered that adsorbed
methoxy may form either CH4 or CH3OH via a proton-transfer
reaction. Our calculations showed that the elementary step to CH4 was
favored by 0.27 eV, on a free-energy basis. Based only on thermodynamics, this would correspond to a molar excess of about 40,000:1 in
favor of CH4, in agreement with the electrochemical experimental
results in which only CH4, and not CH3OH, was observed.
In gas-phase methanol synthesis, the literature suggests that
CH3OH may be formed through a methoxy intermediate.13
However, the electrochemical proton-transfer mechanism from
methoxy to CH4 is not possible in gas-phase synthesis. In electrochemistry, the protons come from solution, and are free to react with
the methyl end of the methoxy (which extends away from the Cu
surface). However, in the gas-phase reaction, the hydrogen addition
would likely come from co-adsorbed hydrogen atoms on the Cu
surface (or from another adsorbate). Co-adsorbed hydrogen would
likely have easier access to the oxygen end of the methoxy (forming
CH3OH), although a kinetic study devoted to methanol synthesis
would be required to distinguish these effects.
The results presented in the current communication will form the
basis of a search for improved catalysts as well as improved understanding of the electroreduction of CO2 on existing materials. The
results suggest that the key enabling step in the formation of
hydrocarbons from CO2 is the protonation of adsorbed CO to form
adsorbed CHO. If adsorbed CHO can be stabilized relative to
adsorbed CO, the necessary overpotential can be significantly
reduced, which will translate directly into a more efficient process.
However, since CO binds only weakly to copper surfaces, materials
that bind CO more weakly will lead to large amounts of gas-phase
CO production. Therefore, materials that bind CHO more strongly,
while binding CO with similar tenacity, may offer the best hope for
future catalyst materials. Many enrichments to the results of the
current study will lead to greater mechanistic insight and will aid in
future materials searches. These possible research directions include
the calculation of an in-depth microkinetic model, the exploration of
different Cu crystal facets, the inclusion of local pH effects, the effect
of local ion concentration, the effect of surface coverage, and the
expansion of the study to metals other than copper.
Acknowledgements
This work was performed as part of the Catalysis for Sustainable
Energy initiative, which is funded by the Danish Ministry of
1314 | Energy Environ. Sci., 2010, 3, 1311–1315
Science, Technology and Innovation. This material is also based
upon work performed as part of the Center on Nanostructuring for
Efficient Energy Conversion (CNEEC) at Stanford University, an
Energy Frontier Research Center funded by the U.S. Department
of Energy, Office of Science, Office of Basic Energy Sciences under
Award Number DE-SC0001060. The Center for Atomic-scale
Materials Design is funded by the Lundbeck Foundation. A.A.P.
acknowledges funding from the Hans Christian Ørsted Postdoc
Programme.
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28 The formaldehyde may competitively desorb from the surface or react
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