The oxidation of water and the reduction of CO2 to fuels
Investigators
Thomas F. Jaramillo, Assistant Professor; Jens K. Nørskov, Professor; Joel Varley,
Post-doctoral Researcher; Lars Grabow, Post-doctoral Researcher, Kendra Kuhl,
Graduate Student; Etosha Cave, Graduate Student; David Abram, Graduate Student;
Desmond Ng, Graduate Student.
Abstract
In this project, we are developing catalysts for CO2 electroreduction and the oxygen
evolution reaction (OER). We report our results in these areas, both experimental and
computational. Three types of catalysts have been investigated in this study: transition
metals, metal sulfides, and metal oxides. Significant advancements have been made both
in terms of catalyst development as well as in providing the knowledge needed for
expediting the development of future catalysts for these reactions.
Introduction
The objective of this project is to develop catalysts for two key energy conversion
reactions: CO2 electroreduction and the oxygen evolution reaction (OER). Developing
catalysts for these reactions can enable technologies that can produce carbon-neutroal
fuels when coupled to renewable energy sources, e.g. wind and solar. Currently, there
are no efficient catalysts for CO2 reduction and the best OER catalysts are scarce precious
metals. In order to produce better electrocatalysts, both experimental and theoretical
approaches were taken. Experimentally, we studied 1) copper and other metal surfaces
for CO2 reduction, 2) metal sulfides for CO2 reduction, 3) polymer-modified Pt surfaces
for CO2 reduction, and 4) manganese oxides as earth abundant and inexpensive OER
catalysts.
Background
There have been several significant advances in CO2 reduction catalysis by our labs
as well as by others in recent times. For instance, at Leiden University in the
Netherlands, research in the lab of Prof. Marc Koper used an on-line mass spectrometer
to understand the mechanism of CO2 reduction on Cu electrodes by comparing the
products of the reduction of suspected reaction intermediates.2 At the University of
Illinois, Prof. Paul Kenis and Richard Masel used ionic liquids as catalysts to
electrochemically convert CO2 to CO at extremely low overpotential.3 And at Stanford
University, theoretical analysis by the group of Prof. Jens Nørskov explored the effect of
different surfaces of Cu on the product distribution and developed activity descriptors for
CO2 reduction metal surfaces to guide the development of better catalysts.4,5 The group
of Prof. Chorkendorff at the Technical University of Denmark worked with the Nørskov
lab to experimentally confirm their prediction that a rougher surface would favor multicarbon products.6 In this report, we describe our most recent advances which are setting a
new standard in the field.
1
Results
CO2 Electrolysis Reactor Design
One difficulty in studying CO2 electroreduction is the need to measure both gas and
liquid phase products. There is no 'standard' method in the field for conducting these
kinds of studies, thus our first goal was to develop a method that would allow for high
sensitivity in identifying and quantifying gaseous and liquid-phase products of reaction.
To achieve this goal we ran electrolysis experiments for 1 hr at a range of potentials using
a custom cell designed to have a large electrode area and small electrolyte volume to
maximize product formation and concentration. CO2 was constantly flowed through the
cell during electrolysis to achieve good mass transport and gas phase products were
measured by a gas chromatograph (GC) downstream of the electrolysis cell. Liquid
phase products were measured by nuclear magnetic resonance (NMR) after electrolysis.
Figure 1 shows a schematic of the cell and examples of the data used to quantify the
products that were formed.
Figure 1. Electrolysis cell used for CO2 reduction with
representative gas chromatograph and NMR data used to
quantify products.
CO2 Electrolysis on Cu Metal
The activity of Cu metal for CO2 reduction is well known, so we began by benchmarking
results obtained using our experimental setup against literature reports. CO2 reduction on
Cu is known to produce methane, ethylene, formate, CO, ethanol, acetaldehyde,
propanol, propionaldehyde, and allyl alcohol along with hydrogen as a side product.7 In
addition to these products, methanol and acetate have been detected in a few studies.8,9
With our newly developed methods, not only did we detect and quantify all of the
previously detected products of CO2 reduction on Cu our system ever reported in the
literature (from all previous studies combined in one single set of experiments), were we
2
also able to detect a number of new, previously unreported products: glyoxal, ethylene
glycol, glycolaldehyde, acetone, and hydroxyacetone. Importantly, we were also able to
track the rates of production for each and every product across a wide potential window.
Figure 2 shows the rates of formation of each product as a function of voltage,
represented in terms of current efficiency, the fraction of current going towards each
product, and in terms of partial current density, which scales linearly with turnover
frequency (TOF).
Figure 2. Current efficiency and Tafel plots of partial
current going to products over the voltage range measured.
Several important observations can be made about the data in Figure 2. At low
overpotential, CO2 is reduced only to the CO and formate, which are believed to have
lower kinetic barriers to formation than the more reduced products. At higher
overpotential, where high kinetic barriers can be overcome, more reduced products
appear. Ethylene and methane are the two major hydrocarbon products. At lower
overpotential slightly more current goes to making ethylene than methane. But at greater
overpotential this trend is reversed. The pathway to multi-carbon products must contain a
C-C coupling step along with proton and electron transfers necessary to all products. As
the overpotential increases, proton and electron transfers become more and more
3
favorable and eventually outcompete the C-C coupling step leading to the production of
less multi-carbon products and more methane and hydrogen.
Another interesting observation is that the partial currents of products with
multiple carbons tend to track each other across the voltage range, suggesting that these
products share common mechanistic pathways. However, the observed multi-carbon
products display a wide range of functional groups; alcohols, carboxylates, ketones, and
double bonds are all present. As the first lab to have identified and quantified this more
complete view of the reaction products and rates of production, we were able to provide
deeper insights into reaction mechanisms that could potentially lead to such diverse
products. An important insight was the realization that the products containing a
carbonyl group have other possible forms. A carbonyl group can react with water to form
a diol or tautomerize to an enol form. By looking at these other forms of the products, we
realized that a plausible mechanism could proceed through the dehydroxylation of the
enol or diol form of the various products. Figure 3 shows this hypothesized pathway in
which the enol or diol form of the products in present as an intermediate on the electrode
surface and either desorbs or is dehydroxylated leading to the formation of more reduced
products.
Figure 3. Proposed pathway for multi-carbon product
formation.
Studying CO2 reduction on Cu allowed us to validate the new experimental methods we
have developed and more importantly, led to several new insights into the reaction
products and pathway. Understanding the mechanism of CO2 reduction has the potential
to allow for control of the product distribution of the reaction and engineer catalysts to
selectively form the desired product depending on the application.
4
CO2 Reduction with Sulfides
Taking inspiration from the enzyme carbon monoxide dehydrogenase (CODH),10 we
explored the CO2 reduction activity of Fe4S4 cubanes supported on highly oriented
pyrolytic graphite (HOPG). While we were successful in attaching the cubanes to the
surface and visualizing them with STM, XPS spectra taken before and after water
exposure showed that the Fe4S4 complexes were unstable (Figure 4A).
Figure 4. A) STM of cubanes attached to an HOPG surface
with XPS showing changes to the Fe and S signal after water
exposure. B) Pyrite film before and after electrochemistry.
We then moved on to exploring thin films of other sulfides. CO2 electrolysis
experiments were performed with thin films of FeS2, FeS, MoS2 and VS2. Again, the FeS materials were unstable (Figure 4B) and hydrogen was the major product detected
during CO2 reduction. MoS2 is a well known catalyst for hydrogen evolution11,12 and
5
100
fficiency
Current E
(%)
produced mostly hydrogen, as did VS2. Figure 5 shows that hydrogen was produced at
high current efficiency at a range of potentials on the sulfides studied.
80
60
40
20
0
-0
1
.2
.3
.4
.R
V vs
.9
.0
.
-1
-1
-1
FeS2
-1
MoS2
-1
FeS
VS2
HE
Figure 5. Current efficiency to hydrogen production on
metal sulfides.
Theoretical studies of CO2 Reduction with Dehydrogenase enzymes
To supplement and guide the experimental efforts on bioinspired systems, we have
explored the relevant reaction mechanisms with first-principles calculations. To this end
we have used the CO dehydrogenase enzyme (CODH) as a model system to study the
redox reaction CO + H2O ↔ CO2 + 2H+ + 2e–.10 The active site of this metalloenzyme
consists of a Ni and Fe4S4 cluster akin to a cubane system, which we have also
investigated. Figure 6 shows a schematic of the studied structures, each of which
provides insight into different mechanisms underlying the catalytic activity of the
enzyme. These include the preferred adsorption sites on the cluster(s) and the role of the
surrounding ligand environment in adsorbate stabilization and selectivity, both of which
can ultimately offer routes to engineering superior catalysts.
6
A
B
C
Figure 6. A) Model of a CODH dehydrogenase enzyme with the immediate ligand
environment as determined by high-resolution X-ray crystallography.13 B) Model of
a CODH enzyme without the ligand environment. C) Model cubane structures of
different Ni-Fe-S stoichiometry. Light green atoms represent Ni, orange represents
Fe, and yellow represents S in all clusters shown.
In Figure 7 we show the relaxed atomic geometries of the adsorbed COOH* and CO*
species and their energetics on one type of model the dehydrogenase species from the
Methanosarcina barkeri bacteria (Mb-CODH).13 We have included the first-order ligand
environment surrounding the active site to explicitly account for stabilization provided by
hydrogen bonding, which we have found has a significant effect in the observed activity.
The free energy diagram describing the CO2 + 2H+ + 2e– ↔ CO + H2O reaction is shown
in Figure 7c, where we find a modest overpotential of -0.47 eV necessary to drive the
reaction electrochemically, in excellent agreement with experiments on other CODH
enzymes.10 We find that the effect of the different cluster atoms also contribute to the
binding energies, as can be see from the COOH* where the C binds to Ni and an O forms
a secondary bond to Fe (Figure 7a). Our future plans include exploiting this as a design
principle for breaking scaling relations between adsorbates that bind via C or O, as
recently detailed in Ref 5.
7
Figure 7. A) Most stable configuration of COOH* and B) CO* adsorbed on the
model CODH dehydrogenase enzyme. C) Gibbs free energy diagram for CO2
reduction to CO* and H2O. The black curve shows no applied potential, while the
red curve shows the minimum potential necessary to make the reaction exergonic.
We have previously shown for a variety of transition metal surfaces that the CO* and
COOH* binding energies both determine the activity in CO2 reduction and follow scaling
relations.14 We have combined our current results on CODH and cubane systems with
ongoing work in a project studying metal surfaces for the same reaction into a
microkinetic model as seen in Figure 8, where we find that both the hydrogen bonding
due to ligands and the “alloyed” surfaces of the CODH and cubanes can lead to
deviations from the transition metal scaling relations. Such deviations can lead to
enhanced activity, as seen by some of the cubane and CODH points that fall closer to the
top of the Volcano Plot in Figure 8. The enhanced activity exhibited by the Mb-CODH
and the Ni2Fe2S4 cubane cluster are being analyzed to further characterize the relative
importance of the geometry (stoichiometry) and stabilization effects of surrounding
ligands.
8
Figure 8. Plot of the Sabatier Volcano for the CO2 reduction activity using the
binding energies of COOH* and CO* as descriptors. The trend line reflects the scaling
relations exhibited by transition metal surfaces, of which only the noble metals are in
the vicinity of the active region. The pink squares represent different Ni-Fe-S cubanes
and the green triangles represent different dehydrogenase enzymes, both of which
show favorable deviations from the scaling relation.
Fischer-Tropsch Synthesis with Nitrogenase enzymes
As was recently discovered experimentally15,16, variants of the nitrogenase enzyme
were found to synthesize a wide variety of hydrocarbons from CO and CN. In addition to
the ammonia synthesis they are known for, this makes nitrogenases extremely promising
catalysts for Fischer-Tropsch synthesis under ambient conditions. The ability to catalyze
such a diverse set of products from simple and abundant reactants makes the nitrogenase
another crucial model system to study for the development of superior bio-inspired
catalysts.
Studies of electrochemical reduction of CO2 on other transition metals
In order to further our understanding of CO2 reduction we studied a variety of
transition metals. The activity Au, Ag, Pt, Zn, Fe, Al, Ti, and Ni plus carbon have been
tested so far and the results are comparable to previous literature reports17 with the
addition of several products that are detectable using our sensitive experimental methods,
many of which have previously gone unreported.
Figure 9a shows the current-voltage data for the tested metals. Clear differences can
be seen in the current generated in a CO2 saturated environment. Figure 9b shows the
percentage of the current that goes to reduce CO2 into any product as opposed to
9
hydrogen. These figures show that although iron and platinum generate high currents at
low overpotential they mostly reduce protons to hydrogen. Although gold, copper and
silver have a moderate total current, they are able to reduce CO2 more effectively than
platinum or iron which generates higher currents overall. Table I, shows the products of
the different metals where red highlights those that are produced on the metal. Currently,
we are conducting deeper investigations into iron, platinum, and gold.
(a)
(b)
(c)
Figure 9: (a) Current-voltage data for 9 transition metals: Cu,
Au, Pt, Ag, Zn, Fe, Al, Ti & Ni plus carbon. (b) Faradaic
Efficiency for CO2 versus Potential for all metals tested plus
carbon. (c) Zoom-in of (b)
Table I. Products of CO2 reduction on metals tested.
The sensitivity of our experimental setup allowed us to detect minor products on a
number of metals for the first time. Figure 10 shows the mechanism of methane
formation on a Cu electrode calculated by DFT.1 A key intermediate along the pathway
is the methoxy intermediate present in step 5. On Cu, it is believed that the majority of
this methoxy intermediate is converted to methane with the next proton transfer in step 6,
leaving behind an oxygen on the surface which is eventually reduced to water. However,
a small amount of methanol is also produced, suggesting that in some cases the reaction
proceeds by protonation of the oxygen of the methoxy group. Assuming a similar
mechanism on other transition metals, Table I shows that Au, the least oxophilic metal
studied, produces methanol but no methane. This can be explained by unfavorability of
step 6 in Figure 10, which would leave a single oxygen atom bound to the Au surface. In
contrast, Zn, Fe, and Ti, all oxophilic metals, produce methane and no methanol,
presumably due to the favorability of the surface oxygen in step 6.
10
Figure 10: Calculated free energy diagram for electro-catalytic CO2 reduction to
CH4 on copper. The black path represents the free energy at U = 0 V. The free
energy at the limiting potential of U = -0.74 V is shown in red.1
Studies of Electrochemical Reduction of CO2 on Pt metal and Polyaniline-Pt
While pure metals can produce a variety of CO2 reduction products, the
overpotential required is larger than desired. Selectivity is also a major concern as a wide
variety of CO2 reduction products are made, in addition to (or instead of) the desirable
hydrocarbons or liquid fuels that could coexist with our existing infrastructure. In an
effort to understand how to impact catalyst activity and selectivity on metal surfaces, we
have begun exploring metal alloys as well as chemically modified surfaces. In one
particular study, we are investigating the impact of polyaniline (PANi) on Pt.
As an aqueous CO2 reduction catalyst, Pt is reported to produce primarily H2 with
0.1% faradaic efficiency for formate18 at 5mA/cm2, though one study does show
methanol, formaldehyde, trace amounts of methane as well.19 Using our recently
developed experimental methods with improved sensitivity limits, we detect and quantify
efficiencies for H2, formate, CO, methanol, and trace amounts of methane on Pt metal in
0.1 M KHCO3 at current densities explored up to 10mA/cm2 as seen in Figure 11.
To investigate effects of polyaniline at the Pt-liquid interface, approximately
10nm of polyaniline was electrochemically deposited on Pt metal foils using cyclic
voltammetry, and these PANi-Pt films were tested using the same CO2 reduction
procedures. While H2 was still the dominant species produced, the product distribution of
among the minor species was significantly altered. The most noticeable changes are an
increase in the production of formate to F.E.’s up to 1.5% vs. 0.3% for pure Pt, and a
reduction in methanol production as seen in Figure 11. The CO production also appears
to increase at higher overpotentials. 13CO2 experiments confirmed that the increased
formate was not attributable to degradation of the PANi film. As formate and CO are
2e- products with the same proposed rate-limiting step, and methanol has a different
proposed rate-limiting step, it is possible that the amine groups of PANi are stabilizing
the intermediate for the 2e- products and destabilizing methanol intermediates. These
effects are critical in impacting the CO2 reduction selectivity, and ongoing efforts are
aimed at understanding how and why the polyaniline can produce such differences.
11
100
1
0.1
0.01
1E-3
Faradaic Efficiency (%)
Current Density (mA/cm
2
)
10
Pt
H2
CO
CH4
HCOOCH3OH
10
1
0.1
0.01
1E-4
-1.0 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4
Potential vs. RHE (V)
-1.0
-0.8
-0.6
Potential vs. RHE (V)
-0.4
PANi-Pt
H2
CO
CH4
HCOOCH3OH
Figure 11: Current densities (left) and faradaic efficiencies
(right) of products produced on Pt and PANi-Pt in 1h
electrolysis in 20sccm CO2-purged 0.1M KHCO3.
Manganese Oxide Water Oxidation Catalysts
A working device for the reduction of CO2 requires a source of protons and
electrons; employing water oxidation, i.e. the oxygen evolution reaction (OER), could
serve as the other half-reaction for this purpose. Through this GCEP-funded effort, we
aimed to develop MnOx catalysts for OER catalysis using two approaches. The first
project involved developing a new and efficient method for synthesizing active supported
MnOx catalysts. This work was based off of work done by Gorlin et al.20 who had
managed to synthesize a highly active thin film of MnOx via anodic electrodeposition
onto a glassy carbon (GC) disk substrate followed by calcination. The ORR and OER
activity of this thin film MnOx is superior to that of the best MnOx catalysts reported in
literature and is comparable to that of the best known precious metal catalysts. However,
the synthesis process is slow, small-scale, and substrate-selective; hence this procedure is
not amenable for commercial use. To overcome the limitations involved yet retain the
excellent ORR activity, we designed a new and efficient synthesis procedure which
incorporated MnOx onto GC particles via a classic impregnation technique. This is
followed by a high-temperature calcination process that generated a nanostructured
morphology for the catalyst. The MnOx-GC particles synthesized displayed comparable
ORR and OER activity with the MnOx thin films as seen in Figure 12(a) and can be
easily loaded onto carbon paper substrate for use in fuel cells and electrolyzers.
The second project involved synthesizing Co-MnOx thin films for OER. MnOx can be
synthesized via cathodic electrodeposition followed by high temperature calcination to
produce an active OER catalyst. To further improve the OER activity, we added a Co
precursor into the deposition solution to synthesize Mn-Co mixed oxides. SEM revealed
that the Mn-Co mixed oxides have a hollow rod-like morphology which is different from
that of the pure Mn oxides and Co oxides. Electrochemical testing showed that the Mn12
Co mixed oxides are more active than the pure Mn and Co oxides as seen in Figure 12(b).
The most active mixed oxide is that with a Co/Mn ratio of 2.2, which corresponds to
MnCo2O4 based on XPS and XRD data. MnCo2O4 requires 470 mV of overpotential to
reach 10 mA/cm2, and stability testing showed that it was still active after 1000 cycles.
(a)
14
Current Density (mA/cm
2
geo
)
12
10
OER
8
6
4
MnOx-GC particles
MnOx thin film (Gorlin et al)
GC particles, calcined
2
0
o
E H2O/O2
-2
ORR
-4
0.1M KOH
5 mV/s, 1600 rpm
-6
0.0
0.2
0.4
0.6
0.8
1.0
1.2
1.4
1.6
1.8
2.0
Potential (V versus RHE)
20
0.64
Mn2O3
15
Overpotential for 10 mA/cm
Current Density (mA/cm
2
geo
)
2
(b)
0.60
0.56
0.52
0.44
10
200 nm
0.48
0
1
2
3
4
5
6
7
OER
Co/Mn Mol Ratio in Film
Mn2O3
Co/Mn 0.95
Co/Mn 2.2
Co/Mn 2.8
Co/Mn 7
Co3O4
5
Eo H2O/O2
0
0.0
0.2
0.4
0.6
0.8
Potential (V vs Ag/AgCl)
Figure 12. (a) Rotating disk voltammograms displaying oxygen electrode
activities of MnOx-GC particles and MnOx thin film. The inset shows an SEM
image of MnOx-GC particles which display a nanostructured morphology. (b)
Cyclic voltammograms showing oxygen evolution activities of Mn2O3, Co3O4,
and Co-MnOx thin films. The inset on the left displays the overpotential
required to reach 10 mA/cm2 for the various catalysts, while the inset on the
right shows an SEM image of the Co-MnOx catalyst (Co/Mn = 2.2) which has a
hollow rod-like morphology.
Progress
We have developed methods for testing the activity of catalysts for CO2 reduction.
Using these methods, a detailed study of Cu catalysis was undertaken and several new
products detected. Based on the product distribution we can hypothesize about likely
reaction pathways and provide insight into reaction mechanisms. On sulfide materials we
detected hydrogen as the major product and found that the some materials studied were
13
unstable at the negative potential required to reduce CO2. Theoretical efforts have been
instrumental in guiding us toward more promising avenues with sulfides. By
understanding how biological catalysts can effectively catalyze CO2 reduction or
ammonia synthesis, we believe that we can utilize those principles to design solid-state
surfaces with similar functionality.
We have also carefully investigated numerous other transition metals. We have
detected a number of new minor products on the transition metals tested so far and can
identify trends in the product distribution based on oxophilicity. Beyond pure metals, we
modified the activity of platinum to produce more formate by the addition of a
conducting polymer to the surface. And towards developing a non-precious metal OER
catalyst, we have worked out new synthesis methods for MnOx catalysts that are suited to
scale-up and explored the addition of Co to improve the water oxidation activity of
MnOx.
Future Plans
Our next steps will focus on furthering our mechanistic understanding of CO2
reduction on Cu and examining the catalytic activity and selectivity of more pure metals
and alloys. Our ultimate goal is to uncover the factors that govern catalytic activity and
selectivity for this reaction, aiming to engineer the surface chemistry appropriately.
Bimetallic surfaces and polymer/metal interactions are promising means to achieve these
goals. In addition to pursuing metallic systems, we will focus on furthering our
mechanistic understanding of CO2 reduction on the nitrogenase and CODH enzymes and
their cubane counterparts. There is much to be learned from the biological catalysts. We
will further characterize the activity and selectivity in terms of adsorbate binding energies
and geometries, looking for favorable deviations from the scaling relations exhibited by
transition metal surfaces. We will also continue studies of the ligand environments in the
real enzymes for inspiration in designing optimal ligand networks that may promote
catalytic activity or selectivity. When coupled with microkinetic modeling, this will give
insight into the design of optimal catalysts.
Publications & Presentations
1.
2.
3.
4.
5.
Kendra P. Kuhl, Etosha R. Cave, David N. Abram and Thomas F. Jaramillo. Energy Environ. Sci.,
2012, Advance Article, DOI: 10.1039/C2EE21234J, Paper
University of California, Santa Barbara, Dept. of Chemistry and Dept. of Chemical Engineering.
"Tailoring Electrocatalyst Materials at the Nano-Scale: Controlling Activity, Selectivity, and
Stability for Energy Conversion Reactions," T. F. Jaramillo, March 2012.
Stanford University, Dept. of Mechanical Engineering, High Temperature Gas Dynamics
Laboratory Seminar Series, Stanford, CA. "Tailoring Electrocatalyst Materials at the Nano-Scale:
Controlling Activity, Selectivity, and Stability for Energy Conversion Reactions," T.F. Jaramillo,
March 2012.
Massachusetts Institute of Technology (MIT), Energy Initiative Seminar Series, Cambridge, MA.
"Tailoring Electrocatalyst Materials at the Nano-Scale: Controlling Activity, Selectivity, and
Stability for Energy Conversion Reactions," T.F. Jaramillo, December 2011.
University of California, Berkeley, Applied Science & Technology, Berkeley, CA. "Catalyzing the
production of clean fuels from renewable energy resources," T.F. Jaramillo, December 2011.
14
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
California Institute of Technology, Chemical Engineering Department Seminar, Pasadena, CA.
“Tailoring Electrocatalyst Materials at the Nano-Scale: Controlling Activity, Selectivity, and
Stability for Energy Conversion Reactions,” T.F. Jaramillo, October 2011.
University of New Mexico, Chemical and Nuclear Engineering Department Seminar,
Albuquerque, NM. “Tailoring Electrocatalyst Materials at the Nano-Scale: Controlling Activity,
Selectivity, and Stability for Energy Conversion Reactions,” T.F. Jaramillo, October 2011.
2011 Annual Meeting of the American Institute of Chemical Engineers (AIChE), Minneapolis,
MN. “Electrocatalytic Conversion of CO2 to Fuels on Metal Surfaces,” T.F. Jaramillo, K.P. Kuhl,
E. Cave, and D.N. Abram, October 2011.
Catalysis for Sustainable Energy (CASE) Seminar Series, Technical University of Denmark,
Lyngby, DK. “Tailoring electrocatalyst materials at the nano-scale: Controlling activity,
selectivity, and stability for energy conversion reactions,” T.F. Jaramillo, August 2011.
Haldor Topsøe Catalysis Forum: Catalysis & Future Energy, Munkerupgaard, DK.
“Semiconductors and catalysts for the production of solar fuels,” T.F. Jaramillo, August 2011.
Lawrence Livermore National Laboratory, Livermore, CA. “Tailoring electrocatalyst materials at
the nano-scale: Controlling activity and selectivity for energy conversion reactions,” T.F.
Jaramillo, July 2011.
Global Climate Energy Project (GCEP) Distinguished Lectureship Series, Exxon Mobil
Corporation, Clinton, NJ. “Tailoring electrocatalyst materials at the nano-scale: Controlling
activity and selectivity for energy conversion reactions,” T.F. Jaramillo, June 2011.
Global Climate Energy Project (GCEP) Distinguished Lectureship Series, General Electric
Corporation, Niskayuna, NY. “Tailoring electrocatalyst materials at the nano-scale: Controlling
activity and selectivity for energy conversion reactions,” T.F. Jaramillo, June 2011.
Global Climate Energy Project (GCEP) Distinguished Lectureship Series, Schlumberger, Inc.
Cambridge, MA. “Tailoring electrocatalyst materials at the nano-scale: Controlling activity and
selectivity for energy conversion reactions,” T.F. Jaramillo, June 2011.
Santa Clara University, Department of Chemistry, Santa Clara, CA. “Nanomaterials for efficient
chemical transformations in energy conversion reactions,” T.F. Jaramillo, May 2011.
Sanford University Energy & Environment Affiliates Program, Stanford, CA. “Nanostructured
Catalysts for Chemical Transformations in Energy,” T.F. Jaramillo, May 2011.
Materials Research Society (MRS) National Meeting, San Francisco, CA. “Electrocatalytic
conversion of CO2 to fuels on metal surfaces,” T.F. Jaramillo, April 2011.
E. Cave, K. Kuhl, and T. F. Jaramillo. Surface Modification of Gold for CO2 Electrochemical
Reduction. 2011 Electrochemical Society Fall Meeting, Oct. 2011, Boston, MA.
K. Kuhl, E. Cave, and T. F. Jaramillo. Electrocatalytic Conversion of CO2 to Fuels on Metal
Surfaces. 2011 Electrochemical Society Fall Meeting, Oct. 2011, Boston, MA.
D. N. Abram, M. Vezie, and T. F. Jaramillo. Conductive Polymer - Metal Nanoparticle
Composites for Electrocatalytic Reduction Reactions. 2011 Electrochemical Society Fall
Meeting, Oct. 2011, Boston, MA.
David N. Abram, Kendra P. Kuhl, Etosha Cave, and Thomas F. Jaramillo. Carbon Dioxide
Electrochemical Reduction on Metals. Fall 2011 GCEP Symposium. Stanford, CA.
Kendra P. Kuhl, Etosha Cave, David N. Abram, and Thomas F. Jaramillo. Carbon Dioxide
Electroreduction on Copper. Fall 2011 GCEP Symposium. Stanford, CA.
Ng J, Baeck S, Jaramillo T. 2011, MnOx Electrocatalyst Development for ORR/OER, GCEP
Seventh Annual Research Symposium
Ng J, Gorlin Y, Jaramillo T, Nanostructured particles of manganese oxide for ORR in alkalinebased fuel cells, in preparation
Cave, E.R, Kuhl, K.P, Abram, D.N. Jaramillo T. F., Electrochemical Reduction of Carbon Dioxide
on Gold, in preparation
Abram, D.N.; Kuhl, K.P.; Cave, E.R.; Jaramillo, T.F. Electrochemical Reduction of CO2 on
Platinum and Polyaniline-Coated Platinum and Modification of Product Distribution, in
preparation
15
References
(1) Peterson, A. A.; Abild-Pedersen, F.; Studt, F.; Rossmeisl, J.; Norskov, J. K.
Energy Env.Sci. 2010, 3, 1311.
(2) Schouten, K. J. P.; Kwon, Y.; van der Ham, C. J. M.; Qin, Z.; Koper, M. T. M.
Chem. Sci. 2011, 2, 1902.
(3) Rosen, B. A.; Salehi-Khojin, A.; Thorson, M. R.; Zhu, W.; Whipple, D. T.; Kenis,
P. J. A.; Masel, R. I. Science 2011, 334, 643.
(4) Durand, W. J.; Peterson, A. A.; Studt, F.; Abild-Pedersen, F.; Norskov, J. K. Surf.
Sci. 2011, 605, 1354.
(5) Peterson, A. A.; Norskov, J. K. J. Phys. Chem. Lett. 2012, 3, 251.
(6) Tang, W.; Peterson, A. A.; Varela, A. S.; Jovanov, Z. P.; Bech, L.; Durand, W. J.;
Dahl, S.; Norskov, J. K.; Chorkendorff, I. Phys. Chem. Chem. Phys. 2012, 14, 76.
(7) Hori, Y.; Murata, A.; Takahashi, R. J. Chem. Soc. Far. Trans. I 1989, 85, 2309.
(8) Hori, Y.; Takahashi, I.; Koga, O.; Hoshi, N. J. Phys. Chem. B 2002, 106, 15.
(9) Le, M.; Ren, M.; Zhang, Z.; Sprunger, P. T.; Kurtz, R. L.; Flake, J. C. J.
Electrochem. Soc. 2011, 158, E45.
(10)
Parkin, A.; Seravalli, J.; Vincent, K. A.; Ragsdale, S. W.; Armstrong, F.
A. J. Am. Chem. Soc. 2007, 129, 10328.
(11)
Chen, Z. B.; Cummins, D.; Reinecke, B. N.; Clark, E.; Sunkara, M. K.;
Jaramillo, T. F. Nano Lett. 2011, 11, 4168.
(12)
Jaramillo, T. F.; Jorgensen, K. P.; Bonde, J.; Nielsen, J. H.; Horch, S.;
Chorkendorff, I. Science 2007, 317, 100.
(13)
Kung, Y.; Doukov, T. I.; Seravalli, J.; Ragsdale, S. W.; Drennan, C. L.
Biochemistry 2009, 48, 7432.
(14)
Abild-Pedersen, F.; Greeley, J.; Studt, F.; Rossmeisl, J.; Munter, T. R.;
Moses, P. G.; Skulason, E.; Bligaard, T.; Norskov, J. K. Physical Review Letters 2007,
99.
(15)
Hu, Y.; Lee, C. C.; Ribbe, M. W. Science 2011, 333, 753.
(16)
Lee, C. C.; Hu, Y.; Ribbe, M. W. Angewandte Chemie-International
Edition 2012, 51, 1947.
(17)
Handbook of Fuel Cells: Fundamentals, Technology and Application;
Hori, Y., Ed.; VHC-Wiley: Chichester, 2003; Vol. 2.
(18)
Hori, Y. W., H.; Tsukamoto, T.; Koga, O. Electrochim. Acta 1994, 39,
1833.
(19)
Brisard, G. M. C., A.P.M.; Nart, F.C.; Iwasita, T. Electrochem. Comm.
2001, 3, 603.
(20)
Gorlin, Y.; Jaramillo, T. F. J. Am. Chem. Soc. 2010, 132, 13612.
Contacts
Thomas F. Jaramillo: jaramillo@stanford.edu
Jens K. Nørskov: norskov@stanford.edu
Etosha R. Cave: cave@stanford.edu
Kendra Kuhl: kpkuhl@stanford.edu
David Abram: dabram@stanford.edu
Desmond Ng: desmond1@stanford.edu
Joel Varley: jvarley@stanford.edu
16