1

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 multi-carbon products.6 In this report, we describe our most recent advances which are setting a new standard in the field.

2

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.

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

Figure 1. Electrolysis cell used for CO2 reduction with representative gas chromatograph and NMR data used to

quantify products.

3

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).

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

Figure 2. Current efficiency and Tafel plots of partial current going to products over the voltage range measured.

4

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.

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.

Figure 3. Proposed pathway for multi-carbon product formation.

5

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).

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 Fe-S 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

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.

6

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.

0

20

40

60

80

100

-1.4-1.3

-1.2-1.1

-1.0-0.9

FeS VS

2

MoS2

FeS2

Cur

rent

Effi

cien

cy (

%)

V vs. RHE

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.

Figure 5. Current efficiency to hydrogen production on metal sulfides.

7

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.

A CB

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.

8

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.

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.

9

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

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.

10

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.

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.

(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.

11

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.

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

12

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 Mn-

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.

-1.0 -0.8 -0.6 -0.4

0.01

0.1

1

10

100

Pt H2 CO CH4 HCOO- CH3OH

PANi-Pt H2 CO CH4 HCOO- CH3OH

Far

adai

c E

ffic

ien

cy (

%)

Potential vs. RHE (V)

-1.0 -0.9 -0.8 -0.7 -0.6 -0.5 -0.41E-4

1E-3

0.01

0.1

1

10

Cu

rren

t D

ensi

ty (

mA

/cm

2 )

Potential vs. RHE (V)

13

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.

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0

-6

-4

-2

0

2

4

6

8

10

12

14(a)

0.1M KOH5 mV/s, 1600 rpm

Cu

rre

nt

De

nsi

ty (

mA

/cm

2g

eo)

Potential (V versus RHE)

OER

ORR

MnOx-GC particles MnOx thin film (Gorlin et al) GC particles, calcined

Eo H2O/O2

0.0 0.2 0.4 0.6 0.8

0

5

10

15

20

200 nm

Mn2O

3

Co/Mn 0.95 Co/Mn 2.2 Co/Mn 2.8 Co/Mn 7 Co

3O

4

Cu

rren

t D

ensi

ty (

mA

/cm

2 geo

)

Potential (V vs Ag/AgCl)

Eo H2O/O2

OER0 1 2 3 4 5 6 7

0.44

0.48

0.52

0.56

0.60

0.64

Co/Mn Mol Ratio in Film

Ove

rpo

ten

tia

l fo

r 10

mA

/cm

2

Mn2O3

(b)

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

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.

14

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. Kendra P. Kuhl, Etosha R. Cave, David N. Abram and Thomas F. Jaramillo. Energy Environ. Sci., 2012, Advance Article, DOI: 10.1039/C2EE21234J, Paper

2. 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.

3. 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.

4. 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.

5. University of California, Berkeley, Applied Science & Technology, Berkeley, CA. "Catalyzing the production of clean fuels from renewable energy resources," T.F. Jaramillo, December 2011.

15

6. 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.

7. 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.

8. 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.

9. 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.

10. Haldor Topsøe Catalysis Forum: Catalysis & Future Energy, Munkerupgaard, DK. “Semiconductors and catalysts for the production of solar fuels,” T.F. Jaramillo, August 2011.

11. 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.

12. 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.

13. 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.

14. 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.

15. Santa Clara University, Department of Chemistry, Santa Clara, CA. “Nanomaterials for efficient chemical transformations in energy conversion reactions,” T.F. Jaramillo, May 2011.

16. Sanford University Energy & Environment Affiliates Program, Stanford, CA. “Nanostructured Catalysts for Chemical Transformations in Energy,” T.F. Jaramillo, May 2011.

17. Materials Research Society (MRS) National Meeting, San Francisco, CA. “Electrocatalytic conversion of CO2 to fuels on metal surfaces,” T.F. Jaramillo, April 2011.

18. 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.

19. 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.

20. 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.

21. David N. Abram, Kendra P. Kuhl, Etosha Cave, and Thomas F. Jaramillo. Carbon Dioxide Electrochemical Reduction on Metals. Fall 2011 GCEP Symposium. Stanford, CA.

22. Kendra P. Kuhl, Etosha Cave, David N. Abram, and Thomas F. Jaramillo. Carbon Dioxide Electroreduction on Copper. Fall 2011 GCEP Symposium. Stanford, CA.

23. Ng J, Baeck S, Jaramillo T. 2011, MnOx Electrocatalyst Development for ORR/OER, GCEP Seventh Annual Research Symposium

24. Ng J, Gorlin Y, Jaramillo T, Nanostructured particles of manganese oxide for ORR in alkaline-based fuel cells, in preparation

25. Cave, E.R, Kuhl, K.P, Abram, D.N. Jaramillo T. F., Electrochemical Reduction of Carbon Dioxide on Gold, in preparation

26. 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

16

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