Electrode Surface Activation and Nanostructuring Effects

ELECTRODE SURFACE ACTIVATION AND NANOSTRUCTURING EFFECTS

ON FUEL CELL PERFORMANCE

A DISSERTATION

SUBMITTED TO

THE DEPARTMENT OF MATERIALS SCIENCE AND ENGINEERING

AND THE COMMITTEE ON GRADUATE STUDIES

OF STANFORD UNIVERSITY

IN PARTIAL FULFILLMENT OF THE REQUIREMENTS

FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

Jason David Komadina

December 2010

This work is licensed under a Creative Commons Attribution- Noncommercial 3.0 United States License. http://creativecommons.org/licenses/by-nc/3.0/us/

This dissertation is online at: http://purl.stanford.edu/zq508hs6390

ii I certify that I have read this dissertation and that, in my opinion, it is fully adequate in scope and quality as a dissertation for the degree of Doctor of Philosophy.

Friedrich Prinz, Primary Adviser

I certify that I have read this dissertation and that, in my opinion, it is fully adequate in scope and quality as a dissertation for the degree of Doctor of Philosophy.

Yi Cui

I certify that I have read this dissertation and that, in my opinion, it is fully adequate in scope and quality as a dissertation for the degree of Doctor of Philosophy.

Paul McIntyre

Approved for the Stanford University Committee on Graduate Studies. Patricia J. Gumport, Vice Provost Graduate Education

This signature page was generated electronically upon submission of this dissertation in electronic format. An original signed hard copy of the signature page is on file in University Archives.

iii Abstract

Fuel cells are an attractive clean energy technology due to the low or zero emissions from operation and the potentially high efficiency. However, several challenges exist that hinder the implementation of fuel cells as a part of portable power devices. Among these challenges are the losses inherent to fuel cell operation. This work presents the results of three studies that attempt to lower activation losses on the fuel cell anode by catalysis or increased reaction area.

The first study focuses on the anode electrochemistry for a novel fuel cell device that utilizes the naturally occurring charge separation in photosynthesis as the electrolyte. The second study investigates the use of an oxide ion conducting electrolyte and a PtRu anode for a direct methanol fuel cell at temperatures much lower than previously considered feasible. The third study examines the fabrication of high-surface area mixed electronic and ionic conducting anodes and their impact on fuel cell performance. The latter two studies are motivated by interest in low-temperature direct methanol fuel cells for mobile devices, while the former was an exploratory investigation into the possibility of “bioelectricity”, which may more appropriately have been called a “photosynthetic fuel cell”.

Photosynthesis energizes electrons obtained through hydrolysis in the thylakoid space and through a series of steps, reduces the charge-carrying protein ferredoxin (Fd). The charge on Fd is used in numerous processes throughout the cell. It is well established that Fd does not readily give up its charge to a bare metal electrode. Therefore, mediators or surface modifiers must be used to capture this high-energy electron. The bioelectricity investigation studied the effects of various chemical modifiers attached to gold electrodes on the electrooxidation of reduced Fd and found that poly-L-lysine covalently bound to a monolayer of mercaptoundecanoic acid on gold resulted in reasonable oxidation kinetics.

Typical solid oxide fuel cells operate at temperatures above 600°C. It was found in the second study that a PtRu anode was effective in low temperature (250-450°C) direct methanol operation using yttria-stabilized zirconia (YSZ), (Y2O3)0.08(ZrO2)0.92, as the electrolyte. In this arrangement, methanol may be used without equimolar quantities of water, in contrast to typical direct methanol fuel cells (DMFCs). Electrochemical impedance spectroscopy (EIS) measurements were analyzed in an effort to better understand the rate-limiting processes.

In the third study presented in this work, the fabrication of high-surface are mixed electronic and ionic conducting (MEIC) anodes is discussed, along with the results of fuel cell characterization with MEIC anodes of high surface area. In particular, fuel cells fabricated using yttrium-doped BaZrO3 (BYZ) electrolytes are studied with high surface area Pd anodes for use with H2 and methanol fuels.

Acknowledgements

I must take a moment to thank all of the wonderful people who have helped me, personally or professionally, along the way. First, I would like to extend my deep thanks to my advisor, Fritz Prinz. Without his constant support and encouragement, this work would not have been possible. Fritz has been as good an advisor as I could ask for, and over the past six years has always been able to motivate me in the lab and remind me how exciting it can be to research. He has a real passion and gift as a mentor, and was always willing to help my progress in any way possible.

My reading committee members, Dr. Paul McIntyre and Dr. Yi Cui, deserve my thanks for their feedback and dedication to reading my work. Along with Dr. Rainer Fasching and Dr. James Swartz, they also engaged in an excellent discussion of my research during my oral defense.

The Nanoscale Prototyping Laboratory (NPL), formerly the Rapid Prototyping Laboratory (RPL), is a large group, and every member past and present has lent a hand in this work. Whether aiding with sample fabrication, testing, or discussing results and theories, the collaborative spirit in this lab is strong and is a phenomenal resource. My coauthors, collaborators, colleagues, and cohorts all deserve my thanks.

Of course, I would not be at this point in my life were it not for the numerous wonderful educators I‟ve had the pleasure to learn from over the years. I have to start as far back as elementary school with my first grade teacher, John Lewis, at Judson Montessori in San Antonio, Texas, who was as happy to teach as I was to learn. Continuing to Lake Country in Minneapolis, I‟ve been lucky enough to encounter dedicated teachers willing to go the extra mile for their students. Through middle school and high school in Edina, Minnesota, there were always those ready to challenge and encourage me in my learning, and the same is true for the instructors in the University of Minnesota Talented Youth Math Program (UMTYMP), which I can‟t recommend enough for young math lovers who have the opportunity to join. I also only have good things to say about my time at Harvey Mudd College and

Stanford University. All of the professors at these institutions are extremely bright individuals with a strong desire to help students produce their best work in both the classroom and the laboratory. Their drive and dedication are an inspiration to me as I finish my doctorate and prepare for my own experience as a professor.

No education is complete without the friendships that develop in the process. I‟ve been fortunate to know a number of wonderful people as friends, many of whom I‟ve also had the pleasure of working with, in our studies or in the lab. Their support during the good weeks and bad weeks has always been appreciated, and it is truly a joy to know all of them.

My heartfelt thanks go to my parents, Kevin and Jayne, and my brother, Greg, too. I don‟t get a chance to see them as often as I would like, but the door is always open, and the table always set. I am sincerely grateful for their, encouragement, love, and faith in my ability to succeed in whatever I set out to do.

I‟ve saved the best for last, as there‟s also a wonderful woman, my fiancée, Rachel. She keeps me fun when I get too serious, and keeps me working when I get too distracted. There is no day so awful that she can‟t fix with one smile, and I could not be happier to have her in my life.

To everyone, thank you.

vii

Table of Contents

Abstract ...... iv Acknowledgements ...... vi List of Tables ...... xi 1 Introduction ...... 1 2 Fuel Cells ...... 5 2.1 General operation...... 5 2.2 Thermodynamic Potential ...... 7 2.3 Reaction kinetics ...... 9 2.4 Losses and efficiency ...... 11 2.4.1 Activation losses ...... 12 2.4.2 Ohmic losses ...... 13 2.4.3 Concentration losses ...... 14 2.4.4 Fuel cell efficiency...... 15 2.5 Fuel cell characterization ...... 17 2.5.1 Current-voltage measurements ...... 18 2.5.2 Cyclic voltammetry ...... 19 2.5.3 Electrochemical Impedance Spectroscopy...... 20 2.6 Summary ...... 22 3 Bioelectricity ...... 24 3.1 Bioelectricity...... 24 3.1.1 Biological Fuel Cells ...... 25 3.1.2 Photosynthesis ...... 27 3.1.3 Challenges ...... 30 3.2 Ferredoxin ...... 32 3.3 Simulations ...... 33 3.4 Experiments ...... 36 3.4.1 Materials and Methods ...... 36 3.4.2 Electrode Surface Modification ...... 37 3.4.2.1 Mercaptopyridine ...... 38 3.4.2.2 Aminoethanethiol ...... 40 3.4.2.3 Poly-L-lysine ...... 41 3.4.3 Analysis of results ...... 44 3.5 Bioelectricity Device Calculations ...... 47 3.5.1 C. reinhardtii photoabsorption model ...... 47

viii

3.5.2 Maximum current density ...... 50 3.5.3 Bioelectricity device fuel cell behavior ...... 51 3.6 Conclusions ...... 52 4 Low Temperature Operation of YSZ DMFC ...... 54 4.1 Direct Methanol Fuel Cells ...... 55 4.2 Experimental ...... 56 4.3 Results and Discussion ...... 59 4.3.1 Electrode Morphology and Composition ...... 59 4.3.2 Fuel Cell Behavior ...... 60 4.3.3 Methanol Cracking ...... 65 4.3.4 Fuel Cell Exchange Current Density ...... 65 4.3.5 Electrochemical Impedance ...... 68 4.4 Conclusions ...... 76 5 High surface area electrodes ...... 77

5.1 Increasing j0 in fuel cells ...... 77 5.2 Mixed electronic ionic conducting (MEIC) electrodes...... 80 5.3 Pd Nanowires ...... 82 5.3.1 Pd Nanowire Growth and Characterization ...... 83 5.3.2 Pd Nanowire Fuel Cell Fabrication Processes ...... 88 5.3.2.1 Dissolution of template on substrate ...... 88 5.3.2.2 Local template etching ...... 89 5.3.2.3 Joining to substrate by electrodeposition ...... 90 5.3.2.4 Thick electroplated Pd substrate ...... 91 5.3.2.5 Drop casting and annealing ...... 91 5.4 Growth of Template on Substrate ...... 92 5.4.1 Aluminum anodization ...... 92 5.4.2 Pd Electroplating in Laboratory-grown AAO ...... 95 5.4.3 Fuel Cell Fabrication with AAO Grown on Substrate ...... 96 5.5 Alternate strategies to high surface area MEIC electrodes ...... 99 5.5.1 Photoresist templating ...... 99 5.5.2 Atomic Layer Deposition of Pd nanotubes ...... 100 5.5.3 Lyotropic phase Pd electrodeposition ...... 102 5.6 Summary ...... 104 6 Conclusions...... 105 References ...... 108

List of Figures

Figure 2.1 – Schematic diagram of H2/O2 fuel cell...... 5 Figure 2.2 – Example fuel cell I-V curve...... 7 Figure 2.3 – Fuel cell losses...... 12 Figure 2.4 – Equivalent circuit model and Nyquist plots...... 21 Figure 3.1 – Bioelectricity schematic diagram and a simplified view of photosynthesis...... 25 Figure 3.2 – Typical microbial fuel cell (MFC) schematic...... 26 Figure 3.3 – Potential energy view of the photosynthetic electron pathway...... 28 Figure 3.4 – Schematic drawing of Chlamydomonas reinhardtii...... 31 Figure 3.5 – Fd-FNR complex and FAD redox...... 33 Figure 3.6 – Structures used for quantum chemical simulations...... 34

Figure 3.7– Structures and pKa values of surface coatings...... 38 Figure 3.8 – PyS-Au results...... 41 Figure 3.9 – PLL-Au results...... 42 Figure 3.10 – PLL-ITO results...... 43 Figure 3.11 – Photoabsorption model geometry...... 48 Figure 3.12 – Estimated bioelectricity device photoabsorption...... 50 Figure 3.13 – Projected fuel cell behavior for hypothetical bioelectricity device...... 52 Figure 4.1 – Schematic diagram of DMFC fuel cell with proton conducting and oxide conducting electrolytes...... 57 Figure 4.2 – SEM image of porous sputtered PtRu anode...... 59 Figure 4.3 – XPS depth profiling of PtRu anode after testing...... 61 Figure 4.4 – Current-voltage and power density data at 250°C...... 62 Figure 4.5 – Current-voltage and power density data at 350°C...... 64 Figure 4.6 – Peak power density (mW/cm2) plotted on a log scale...... 64 Figure 4.7 – Kinetic parameters determined by Tafel fitting...... 67 Figure 4.8 – Sample EIS data taken at 300°C...... 69 Figure 4.9 – Sample EIS data taken at 250°C...... 70 Figure 4.10 – Typical results from fitting the EIS data to four R||C loops...... 72 Figure 4.11 – Area specific capacitances of four R||C loop model...... 73

Figure 5.1 – Example activation losses for different values of j0...... 78 Figure 5.2 – Fuel cell behavior for corrugated thin film structure...... 79 Figure 5.3 – Schematic diagram of SOFC with one structured electrode...... 80

Figure 5.4 – Schematic of close-packed nanowire geometry...... 83 Figure 5.5 – TEPC-templated nanotube-like structures...... 84 Figure 5.6 – AAO template Pd nanowires...... 85 Figure 5.7 – Comparison of AAO-templated Pd wire rinsing methods...... 85 Figure 5.8 – Characteristic CVs of Pd film and Pd film with wires...... 86 Figure 5.9 – TEM analysis of Pd nanowire...... 87 Figure 5.10 – Schematic diagram of a BYZ-based fuel cell built on locally etched AAO membrane...... 89 Figure 5.11 – Thin-film supported nanowires joined to substrate by electroplating...... 90 Figure 5.12 – SEM image of commercial AAO with 6.5 m of Pd deposited by electroplating...... 91 Figure 5.13 – Current-time behavior of anodization...... 94 Figure 5.14 – TEM image of AAO on Pd supported by Si wafer...... 94 Figure 5.15 – Pd deposition in lab-grown AAO pore channels...... 95 Figure 5.16 – Free standing AAO/BYZ membrane on a silicon wafer piece...... 96 Figure 5.17 – Proposed fuel cell fabrication process for Pd nanowire anode SOFCs...... 98 Figure 5.18 – Pd pads electrodeposited through a photoresist shadow mask...... 99 Figure 5.19 – ALD Pd deposited in AAO pore channels for SOFC...... 101

Figure 5.20 – CVs of flat and lyotropic H1 phase electroplated Pd surfaces...... 103

Figure 5.21 – Fuel cell behavior of SOFC with lyotropic H1 phase deposited Pd anode...... 103

List of Tables

Table 2.1 – Common fuel cell types and key design parameters ...... 6

Table 2.2 – Typical j0 values for selected catalysts...... 11 Table 3.1 – Summary of CV results and analysis...... 45 Table 4.1 – EIS loop activation energies (eV)...... 73

1 Introduction

In 2007, the world energy demand, as measured by the energy actually sold on the market, was about 495 quadrillion Btu (“quads”) (1). This is not a very useful number to most people. We can convert this amount into watts and find that this is about 16.6 TW, on average over the year. People tend to think in terms of things we can comprehend. For the average citizen of the world, 16.6 TW is about the same as if every last person on the planet was running 40 light bulbs at 60W, day and night all year. Not every household or family – every person. This perspective still doesn‟t tell the story very well. The fact of the matter is that there‟s not a good way to say how much energy is fair for a given person to use. Or how much energy the world should ideally be consuming. The fact is that the energy demand is rising, and has been rising for some time, and will continue to rise (by a projected 10 quads annually (1)), as our population grows and quality of life improves. The heart of the issue, and what really matters, is where the energy comes from. There may be room to argue that some new technologies bring less to quality of life than they do to the energy demand table, but improving how we get – and use – energy, is an easier pill to swallow.

The United States consumes, on average, 20% of the world energy, with only about 5% of the world population. The U.S. also produces around less energy than it consumes (94.6 quads versus 73 quads produced in 2009). The numbers are further skewed when considering consumption of oil produced and consumed by the U.S.: about 8,500 barrels produced and about 19,500 barrels consumed daily. Again, the point is not the overall size of the numbers. The issue here is that the U.S. uses more energy than it produces, and is the largest consumer of energy in the market. China is poised to become the bigger customer, but not per capita, at least not yet.

There are a number of arguments about renewable energy and the environment, national security, the economy, and so on. These are beyond the scope of this dissertation. I have one point about why we should care about renewable energy: efficiency is money. I don‟t mean “drive a hybrid car, get rich”. I mean that by more

1 efficiently using the limited resources available to us, the more we can do with those resources. The Sun is often used as the zenith of renewable energy, and with good reason. It is, at present, a practically infinite reservoir pumping orders of magnitude more energy our way than we can currently use (about 120,000 TW, not counting what‟s reflected by the atmosphere and the surface (2)). Granted, the more light we absorb and convert into electricity, the less is absorbed by the Earth‟s crust or reflected into space. One might argue that this has fewer, and far less obvious, consequences to our home planet than mining for coal and drilling for oil, both of which are limited resources. We know these resources are limited because we know how big the Earth is and how much of it is taken up by things that are not oil. The short version of my personal argument for renewable energy technologies: the greater the portion of our demand that is met by renewable technologies, the slower we use the limited, but energy-rich and (increasingly less) easily accessible reservoir.

The good news for those that acknowledge the evidence for renewable energy technologies is that the percentage of demand met by such technologies is slowly but surely increasing. In 2009, the United States consumption included about 8% from hydroelectric power, wood, biofuel and biomass, wind, geothermal, and solar sources. The United States production from these sources was closer to 11% (1). Not surprisingly, these percentages represent about the same amount of energy.

Often when developing engineering solutions to a problem, it can be extremely beneficial to turn to the natural world for inspiration. Plants absorb about 90TW of sunlight (2) to gain the energy, through photosynthesis, needed to survive. We have developed a similar technology in photovoltaics and can convert sunlight directly into electrical energy, or thermal energy in the case of solar water heaters. There is one problem with the Sun, though – sometimes we can‟t see it, and we use energy around the clock. Plants solve this problem by storing the light energy gained during the day as chemical energy – in the form of ATP, NADPH, leaves, potatoes, etc. Thus we see in photosynthesis a system that uses and stores energy from light as chemical energy.

And we see efforts to convert sunlight into electrical energy as well as store it as chemical energy using, for example, solar-charged batteries.

One proposed alternative strategy for storage of energy from cyclic power sources like solar and wind is the “hydrogen economy”, in which a portion, or all, of the energy converted is stored as chemical energy in the form of molecular hydrogen. The hydrogen can then be transported and used on-demand, similar to batteries or flow batteries charged by a solar or wind farm. Again we can draw the parallel to photosynthesis; plants convert sunlight during the daylight hours to store energy as sugar during night. In observing natural solutions such as this to engineering problems (how to convert energy and store some for future use), we gain the benefit of millennia of evolution and natural selection, which can be viewed as a sort of optimization process for biology.

In the first study discussed, a proposed device interfaces manmade nano-sized electroactive probes with photosynthesis in vivo. By inserting two electrodes into the system, an electrochemical circuit is completed, and a device with similarities to both solar cells (light absorption and photosynthetic action) and fuel cells (complementary oxidation and reduction reactions) is made. Some work has been done with live cells by other members of our laboratory, however my work focuses on an in vitro study of the interaction between one of the key proteins in photosynthesis, ferredoxin, and one of the electrodes. To improve the characteristics of this interaction, we turned to the reaction mechanism found in the cells, and found that by treating our electrodes with surface adsorbed molecules to present a more familiar reaction site, ferredoxin is more likely to donate an electron to a metal surface.

In the final study presented, we again find inspiration in nature. Every living cell is tasked with producing a large number of chemical compounds and proteins for use in and out of the cell, yet cursed with a limited space in which to do so. Within each cell, are mitochondria, which nearly everyone who has had biology in grade school knows as the “power plant” of a cell: small organelles tasked with producing most of the chemically energetic adenosine triphosphate (ATP) used throughout the cell (3). To

3 get the most reaction in a tight space, the inner membrane is folded to create a high surface area within a fixed volume. This led us to study high surface area structures for fuel cell electrodes. Previous research in our lab has shown that folding (“corrugating”) a fuel cell membrane can increase the total current density and power density for a fixed device size. The research presented here is an effort to study the effects of a high surface area interface between the electrode and fuel. The hypothesis is that an increased surface area at this interface can help reduce so-called “activation losses” by increasing the amount of reaction per time without changing the rate of reaction per unit reaction site area.

In all of the studies discussed in the following chapters, the end goal is new and/or improved devices that convert chemical or radiation (light) energy directly into electrical energy. In doing so, we hope to add to the state-of-the-art for technologies that reduce the role of limited-supply resources and lower-efficiency energy conversion processes such as combustion.

2 Fuel Cells

This chapter discusses fuel cells in general terms in order to provide sufficient background for the work presented in subsequent chapters. The basics of fuel cell operation (4) are explained, followed by discussions of fuel cell thermodynamics; reaction kinetics; relevant losses and fuel cell efficiency, including comparison to combustion engines; and the characterization of fuel cells, including impedance spectroscopy analysis.

Figure 2.1 – Schematic diagram of H2/O2 fuel cell. + - + - Anode reaction is H2 → 2H + 2e . Cathode reaction is 2H + 2e + ½ O2 → H2O. Thus the net reaction is H2 + ½ O2 → H2O, 1.23 V under standard conditions.

2.1 General operation

Fuel cells are electrochemical devices in which ions and electrons are separated by chemical reactions, and transported from one electrode to another by different pathways – the electrons are transported through an external load, and ions through an electrolyte. The operation of a hydrogen fuel cell is illustrated in Figure 2.1. On the anode (oxidizing electrode) side, hydrogen is split into protons and electrons. This is known as the hydrogen oxidation reaction (HOR). The electrons pass through the load and the protons are transported through the electrolyte, which is electrically insulating. The protons recombine with the electrons on the cathode (reducing electrode) side,

5 where oxygen is reduced to form water. This is the oxygen reduction reaction (ORR). The net reaction is the combustion of hydrogen. The process is similar for other fuel cell types, but may involve conduction of different ions or oxidation of different fuels. In most fuel cells, the net reaction is a combustion reaction. What principally differentiates the fuel cell arrangement from a hydrogen combustion engine is the separation of the reaction into the two electrochemical half reactions, HOR and ORR.

During normal operation of a H2/O2 fuel cell, the separation of fuel and oxidant (and thus HOR and ORR) results in the development of a concentration gradient of H+, as well as a voltage gradient. The voltage gradient is lowered by the net reaction proceeding forward; the greater the current, the lower the cell potential. It is more useful to consider the area-specific current density and power density than current and power because the intrinsic information facilitates comparison of different fuel cells.

There are a number of arrangements of fuel cells. Hydrogen gas is the most common fuel, but depending on the electrolyte material, the charge carrying ion and ideal operating temperature can be different from the example H2/O2 fuel cell discussed in this chapter. Table 2.1 lists a few examples of other fuel cell systems.

Table 2.1 – Common fuel cell types and key design parameters Type Electrolyte (Ion) Temp. Range Fuels Catalysts

Polymer Electrolyte Solid acid polymer H2 Pt-group + ~ 100°C Membrane (PEMFC) e.g.: Nafion (H ) CH3OH

Ceramic H Solid Oxide Fuel Cell 2 e.g.: BYZ (H+) > 600°C CH Ni-cermets (SOFC) 4 YSZ (O2-) CO

Phosphoric Acid Fuel Phosphoric Acid ~ 200°C H Pt-group Cell (PAFC) (H+) 2

Alkaline Fuel Cell Potassium < 200°C H Pt-group (AFC) Hydroxide (OH-) 2

Figure 2.2 – Example fuel cell I-V curve. Region I – Activation losses dominate, j ~ exp(V). Region II – ohmic losses dominate, linear shape. Region III – Mass transport losses appear and voltage drops as limiting current is approached.

An example of fuel cell current-voltage behavior is shown in Figure 2.2. The figure indicates the three regimes of operation as defined by the limiting factor to performance. The remainder of this chapter will discuss the origins of fuel cell parameters and the characterization of fuel cells, beginning with the thermodynamics of steady-state operation.

2.2 Thermodynamic Potential

The thermodynamic potential of a fuel cell (or open circuit voltage, OCV) is determined by the thermodynamics of the two half reactions. For the example of a

H2/O2 fuel cell (proton conducting electrolyte), the reactions are

+ – Anode: H2 2H + 2e 2.1 + – Cathode: ½O2 + 2H + 2e H2O 2.2

Net: H2 + ½O2 H2O 2.3 The standard free energy change of the net reaction is calculated from the standard formation energies of the species involved:

GO GO GO GO rxn f, H2O f, H2 f, O2 2.4 It is uncommon for fuel cells to operate under standard conditions, and therefore, deviations from standard state must be accounted for. The nonstandard reaction free energy change is

O v Grxn (T, P) Grxn RT ln ai 2.5 v Where R is the ideal gas constant, T is the absolute temperature, and ai is the activity of species i raised to the power of its stoichiometric coefficient in the reaction. For the

H2/O2 fuel cell net reaction (Equation 2.3), the activity product in Equation 2.5 is

a v H2O 1 ai a a1/ 2 p p1/ 2 2.6 H2 O2 H2 O2 Gibbs free energy change can be related to an electrostatic potential difference through the factor –nF, where n is the number of electrons transferred in a reaction and F is Faraday‟s constant. The resulting expression is the Nernst equation:

RT E E O ln av nF i 2.7

The standard reaction potential, E°, for the H2/O2 fuel cell is 1.23V. The reaction in Equation 2.1 is defined to have an equilibrium potential of 0 V, and when it occurs at a Pt surface, the system is known as the normal hydrogen electrode (NHE, equivalent to “standard hydrogen electrode”, SHE), and is often used as a reference for other reaction potentials. Since the reaction potential of the fuel cell is equivalent to the potential difference between the two electrodes, the H2/O2 fuel cell cathode reaction therefore has a standard equilibrium potential of +1.23 V.

For gaseous reactants and products, the activity is typically taken to be the partial pressure of the species. This does assume some ideality, but is a reasonable approximation for predicting cell voltage. It is also important to account for entropic considerations with deviation from standard temperature. Thus a more complete form of the Nernst equation is given by

O S(T T0 ) RT v E E ln ai 2.8 nF nF

The entropic term is negative for H2/O2 fuel cells, and thus the equilibrium potential has a tendency to decrease with increasing temperature. For some fuel cells, such as direct methanol fuel cells (DMFCs), the entropic term is positive, resulting in an increasing equilibrium potential with increasing temperature. The parameter T0 is the temperature used to determine the standard potential.

The Nernst equation is a valuable tool in understanding how the OCV ( E of reaction) of a fuel cell varies with temperature and species activity. Increasing reactant pressures has only a logarithmic effect on the OCV, but can still have a significant impact on the overall fuel cell performance by increasing the current and power densities at a given cell voltage, as evidenced by the concentration dependency of the kinetics of reaction.

2.3 Reaction kinetics

The current density from an electrochemical reaction is directly relatable to the rate of reaction in terms of molar flux:

j nFJ 2.9 where J is the number of moles reacted per unit area per unit time. The number of moles reacted per unit time is related to the density of reaction sites, probability of excitation, and probability of reaction completion upon excitation. If we consider the electrochemical reaction O + e- ↔ R, then the net rate of reaction is the sum of the forward and reverse reaction fluxes and is expressable as

Ea,for Ea,rev J net J for J rev cO ffor exp cR f rev exp 2.10 kBT kBT * where c i is the surface concentration of species i, f is the probability of an activated state transitioning to the product state divided by the activated state lifetime, Ea is an activation energy, and kB is the Boltzmann constant. Rewritten as current density, the above equation becomes

Ea,for Ea,rev jnet nFcO ffor exp nFcR f rev exp 2.11 kBT kBT At equilibrium, net current density is zero, and therefore the forward and reverse current densities must be equal. The magnitude of each at equilibrium is reffered to as the “exchange current density”, j0. If electrical potential effects are included, then deviation from equilibrium can be described by the Butler-Volmer equation:

F (1 )F j j exp exp net 0 act RT act RT 2.12 where act is defined as E – V, the deviation from OCV towards lower potential and is the “transfer coefficient” of the reaction, and represents the relative impact of each reaction activation energy due to an electrical potential across the reaction interface. The transfer coefficient is necessarily between 0 and 1, with a value of 0.5 indicating that the forward and reverse reaction barriers are affected equally by electrical potentials.

It is important to note that the Butler-Volmer equation describes the current-voltage relationship for a single step, single electron reaction, and there are at least two such reactions in any fuel cell. For multistep reactions or reactions involving multiple electrons, it is common to assume that the reaction kinetics can be described by a single limiting step. In a fuel cell, unless specific information is known about the potential drop at the oxidant/cathode interface and the anode/fuel interface, it must be assumed that a single reaction step on either the anode side or cathode side limits the whole fuel cell current density. This is because the net current density of the anode reaction must equal the net current density of the cathode reaction in steady-state. For

H2/O2 fuel cells, it is reasonable to assume that the ORR on the cathode is limiting. One can form an idea for which reaction will limit fuel cell performance by comparing exchange current density for the reactions involved. Table 2.2 lists some typical values for j0. The orders-of-magnitude difference between HOR and ORR seen in the table supports the assumption that the ORR limits behavior.

Table 2.2 – Typical j0 values for selected catalysts. Given to the nearest order of magnitude, for H2/O2 fuel cells with proton conducting acidic polymer electrolyte at 300K and 1 atm. -2 Reaction Catalyst j0 / A-cm Pt 10-3 HOR Pd 10-4 Au 10-6 Pt 10-9 ORR Pd 10-10 Au 10-11

It is also important to discuss the concept of the “triple phase boundary” (TPB). Equations 2.1 and 2.2 show the anode and cathode reactions for a hydrogen/oxygen fuel cell. In each reaction, there exists a gas phase component, electrons, and ions. In order for the reaction to occur, all three phases must be present. A dense coating of Pt on the electrolyte surface will prevent fuel cell operation because it blocks ion transport. A porous Pt coating, however, can act as the current collector and catalyst, while allowing for interaction of the gas, electronic, and ionic components.

2.4 Losses and efficiency

As seen in Figure 2.2, the cell voltage deviates further from the OCV with increasing current. This voltage loss can be considered part of the energy cost of fuel cell operation. The three most significant sources of loss are the activation losses from reaction kinetics, act as in §2.3, ohmic losses from charge transport, ohm, and mass transport (“concentration”) losses, conc. The overall I-V behavior of a fuel cell is determined by the sum of these losses:

Vcell OCV act ohm conc 2.13 Figure 2.3 illustrates how each of these losses contributes to the overall I-V behavior of a fuel cell. The origins of the losses are presented below.

Figure 2.3 – Fuel cell losses.

I-V behavior is dictated by the sum of losses from activation, act, charge transport, ohm, and mass transport, conc.

2.4.1 Activation losses

As discussed in §2.3, the reaction current density is governed by the exchange current density (reaction kinetics) and the applied potential, in the form of act. This is the activation loss, and act represents deviation from the OCV. The activation loss can be quantified in terms of the kinetic parameters by rearranging the Butler-Volmer equation using one of two approximations for different operating conditions.

Linearization, assuming j << j0 (low act)

F (1 )F j j 1 1 net 0 act RT act RT 2.14 RT j act 2.15 F j0

Tafel approximation, j >> j0 (high act)

F j j exp 0 net 0 act RT 2.16 RT j act ln 2.17 F j0 In the first (low overpotential) approximation, the Equation 2.12 has been linearized, allowing us to express the activation loss as a linear function on current density and j0. In the second (high overpotential) approximation, the reverse reaction current is assumed to be negligible, and only the forward reaction branch is considered, resulting in a logarithmic expression for activation loss. The most important observation from these expressions is that a higher j0 results in lower activation loss. There are two ways to compare a fuel cell with high j0 versus one with low j0: at a given current density, less potential is lost; alternatively, for a given activation loss, a higher current density is achieved. The exponential dependence of the net current density on activation overpotential can be seen in Figure 2.3.

It is important to consider both the cathode and anode losses from activation. However, as mentioned in §2.3, the ORR tends to be sluggish relative to the HOR, and one can often assume that the anode contributions to activation overpotential are negligible in comparison to the cathode contributions for a H2/O2 fuel cell.

2.4.2 Ohmic losses

Charge transport results in a linear loss in potential with increasing current as described by Ohm‟s Law. For fuel cells, there are two sources of charge transport (“ohmic”) loss. The first is from the transport of electrical current between the fuel cell and load. The second is due to ion transport in the electrolyte. The total loss from charge transport is thus given by

Lelec Lion ohm i(Relec Rion ) i 2.18 Aelec elec Aion ion

13 where i is the absolute current, and L and A are the path length and area of conduction, respectively. The subscripts indicate electronic or ionic conduction. Because the conductivity of most electrical conductors is several orders of magnitude greater than the ionic conductivity of even the best electrolytes, electrical conduction losses are often negligible. The area of ionic conduction, Aion, is often the cell area used to determine current density. In this case, the ohmic loss can be rewritten as

Lion ohm i(Relec Rion ) j j ASRion 2.19 ion The quantities j and L/ (area-specific resistance, ASR) are both intrinsic quantities, allowing for comparison between two fuel cells. Fuel cell ohmic loss manifests as a linear regime of I-V behavior, as seen in Figure 2.3. Lower ASR results in higher cell performance. Therefore, it is important to select an electrolyte that has a high ionic conductivity. It is critical to realize that the best conductor is not always the best choice for the electrolyte, due to operating temperature, fuel and catalyst permeability, and other engineering considerations.

2.4.3 Concentration losses

The last major source of loss occurs at very high current density (low cell potential), when the design of a fuel cell assembly is such that mass transport is the limiting factor. That is, reaction rates are so high that fuel or oxidant (or both) cannot be delivered sufficiently quickly to maintain higher reaction rate. Mass transport loss is present throughout the I-V behavior, although typically negligible until the limiting current density, jL, is approached, as seen in Figure 2.3. The relation between concentration loss and current density is given by

jL conc C ln 2.20 jL j where the factor C is a constant that can be determined empirically for a given fuel cell, and is sometimes approximated as

RT 1 C 1 F 2.21 The form in the above expression is attributed to the effect that depleted fuel concentration at the reaction interface has on the Nernst potential and on reaction kinetics.

The limiting current density is a product of fuel delivery and electrode design and is typically calculated as

eff o jL nFD cR 2.22 eff where D is the effective diffusivity of reactant, c°R is the bulk reactant concentration far from the reaction surface, and is the diffusion length (often assumed to be the electrode thickness in the case of porous electrodes that also serve as a “gas diffusion layer”). Therefore, the limiting concentration losses of a fuel cell may be abated by improving the diffusion characteristics (for example, by minimizing the tortuosity of a porous electrode or gas diffusion layer, or using a different carrier gas). Alternatively, the bulk concentration of the reactant may be increased (for example, by increasing the reactant flow rate), although this has implications in the overall efficiency of the device.

Both anode and cathode concentration losses can affect fuel cell performance. Like activation losses, mass transport on the cathode may be much slower than that on the anode. In such a case, the anode mass transport losses may be considered negligible.

2.4.4 Fuel cell efficiency

The total efficiency of a fuel cell is the product of the thermodynamic efficiency, the voltage efficiency, and the fuel efficiency:

net thermo voltage fuel 2.23 The thermodynamic efficiency is ratio of the Gibbs free energy of reaction (useful energy) to the enthalpy of reaction (maximum heat energy). For a H2/O2 fuel cell, the higher heating value of water is used, as the formation of liquid water constitutes a

15 loss when the fuel cell operates at a temperature above the boiling point. The expression is thus

Grxn thermo 2.24 H HHV The thermodynamic efficiency is a property of the fuel cell chemistry, and is, like the thermodynamic potential, only minimally affected by changes in reactant activities.

For H2/O2 fuel cells near 100°C, the thermodynamic efficiency is approximately 83%.

The voltage efficiency is the result of deviation from thermodynamic equilibrium and is determined by

Vcell voltage (1 act ohm conc ) / Ethermo 2.25 Ethermo It is clear from the above expression that the fuel cell losses have a direct impact on the efficiency of a fuel cell.

Lastly, one must consider the efficiency of the fuel delivery. This is typically represented by a “fuel utilization” coefficient, . A factor of = 1 indicates that all of the reactant delivered to the fuel cell is consumed. For air-breathing fuel cells (i.e. those with the cathode open to air at atmospheric pressure), it is only necessary to consider the anode side fuel consumption. However, if the oxidant is delivered as compressed air or oxygen, for example, utilization at both anode and cathode must be considered. Fuel utilization, and therefore overall efficiency, is dependent on fuel delivery design. There are two approaches to fuel delivery: stoichiometric delivery, which delivers fuel proportional to the rate of consumption; and constant flow rate, which delivers sufficient fuel for the highest current densities required at all operating conditions. In the former case, the fuel efficiency can be assumed to be roughly constant. In the latter case, the fuel efficiency increases with current density.

The net efficiency of a fuel cell can now be written as

Grxn Vcell net 2.26 H HHV Ethermo The net efficiency follows the shape of the I-V curve for constant stoichiometry fuel delivery. For constant flow fuel delivery, the net efficiency tends to increase as current density increases as increases, but then falls off as the voltage efficiency becomes very low.

With well-designed fuel delivery systems, fuel cells can achieve very high efficiencies, near the thermodynamic limit (~80%) at low current densities, and often near 40-50% at peak power density. In contrast with fuel cells, in which chemical energy is converted directly into electrical energy, the thermodynamic efficiency of internal combustion (IC) engines is limited by the Carnot cycle. The maximum theoretical efficiency of an IC engine is typically less than 50%, with increasing efficiency with size and operating (“rejection”) temperature. However, the experiments discussed in subsequent chapters utilize a constant flow rate to eliminate significant concentration losses from the analysis of the fuel cells tested, and therefore suffer from extremely poor net efficiency. As an example, a common flow rate used in our experiments is 30 sccm H2. By converting volume to moles gas and multiplying by nF (n = 2 for H2), the maximum current at this fuel delivery rate is about 3.9 A. Research devices in our lab tend to have an absolute current output on the order of mA at best, resulting in net efficiencies lower than 1%.

2.5 Fuel cell characterization

The physics and chemistry of fuel cell operation is useful to predict performance, but it is difficult to account for all of the variations from ideality. Therefore, characterization of fuel cells is important. Using the information from measurements in combination with the theory behind the behavior, problems can be identified and improvements can be made. This section will discuss three common electrochemical characterization tests, the I-V measurement, cyclic voltammetry (CV), and electrochemical impedance spectroscopy (EIS). These methods are all discussed in more detail in the literature (4; 5).

2.5.1 Current-voltage measurements

The current voltage behavior of a fuel cell represents the steady-state operation of the device. The most commonly employed methods for measuring the I-V behavior are potentiostatic current measurements and galvanostatic voltage measurements. In potentiostatic measurements, the voltage is held constant and the current is measured over a sufficiently long period of time to ensure steady state has been attained. The potential is then stepped to the next point in the measurement. Galvanostatic measurements step current and measure voltage. An alternative to this method is to use a linear sweep of current or voltage. Linear sweep methods are faster, but risk distorting the I-V behavior by not allowing the system to reach steady state. Therefore, when linear sweep methods are employed, it is critical that the scan rate is slow enough to accurately measure the I-V behavior. Typically mass transport effects are most severely distorted in a linear sweep since gas diffusion is the slowest process in the fuel cell. Activation and ohmic effects respond much faster and are less impacted by the scan rate.

I-V curves give an overall impression of fuel cell performance and can be used to calculate the power density output of the cell, as well as obtain estimates for the ohmic resistance from ion transport, and in some cases, make estimates of the limiting reaction exchange current density and transfer coefficient. The power density is obtained from the product of the cell voltage and current density. For I-V curves with a distinct linear regime, as in Figure 2.2, the ohmic resistance can be obtained from the slope of the linear portion of the curve. To obtain kinetic information, the I-V curve must be corrected for ohmic loss, and then a linear fit of the natural logarithm of current density versus overpotential is made. This is called Tafel fitting because it uses the Tafel approximation of activation losses to construct a plot of ln j vs act. The slope of the linear fit can be related to the transfer coefficient, and the axis intercept at zero overpotential is ln j0.

2.5.2 Cyclic voltammetry

Cyclic voltammetry (CV) is an electrochemical analysis method that combines two linear voltage sweeps. The working electrode potential is varied linearly at a fixed rate between two endpoint voltages relative to a reference electrode (such as NHE or AgCl/Ag). The resulting current measurement can be analyzed to obtain various thermodynamic and kinetic parameters from the curve shape.

The governing equations for current in a linear potential sweep require balancing diffusion considerations with kinetics of reaction. The measured current increases exponentially as the potential approaches and passes the equilibrium potential for a reaction between the working electrode and electrolyte system being studied. For fixed electrodes, the current will peak and subsequently fall with the square root of time due to depletion of reactants at the electrode surface and diffusion limitation. If the reaction is reversible, the complimentary redox reaction is observed during the return voltage sweep, with the opposite sign in the current measurement. The separation between oxidation and reduction peaks and the peak height are determined by the diffusivity and bulk concentration of reactants, sample area, voltage scan rate, and reaction kinetics. Therefore, it is possible to determine some of these parameters given sufficient information about the system. For example, in §3.4, CV measurements are used to extract kinetic information about the reaction between a gold electrode and a protein in solution, and in §5, CV is used to determine the electrochemically active surface area enhancement of nanostructured electrodes relative to a control sample. Further detail of the specific CV analyses performed is presented with the relevant experimental results in subsequent chapters.

The electrode geometry (i.e. diffusion character) also plays a role in determining the curve shape. It is also common to use a rotating electrode or to stir the solution to help eliminate diffusion considerations. In the experiments presented in this dissertation, a fixed planar electrode is used for all CV experiments.

2.5.3 Electrochemical Impedance Spectroscopy

Electrochemical impedance spectroscopy (EIS) is a powerful analysis tool that can be used to observe the various losses in a fuel cell. In “potentiostatic” mode, the input signal is a voltage of the form v(t) = vDC + vAC sin it, and the resulting current is measured as a function of time to calculate the frequency-dependent cell impedance.

The frequency of the AC component, i, is stepped through a wide bandwidth. It is useful to obtain current response measurements for the range of frequencies as well as for several values of vDC. The “galvanostatic” mode is similar, with current signal input and voltage measured.

Impedance is the AC analog of ohmic resistance, and is calculated as Z( = V( /I( , yielding a magnitude, |Z|, and a phase, :

Z( ) Z exp(i t) Z cos( t) i Z sin( t) ZRe iZIm 2.27 where i is the imaginary unit. The results of EIS measurements are typically plotted one of two ways. Bode plots display the phase and the log of the magnitude of the impedance versus frequency on a log scale. Nyquist plots display –ZIm versus ZRe. The discussions of EIS in this work are based on the interpretation of Nyquist plots by fitting to equivalent circuit models.

Electrochemical reactions can be modeled as a resistor (kinetic impediment to reaction) parallel to a capacitor (electrochemical double layer). An ideal R||C component in the model manifests as a semicircle in the Nyquist plot. The diameter of the semicircle is equivalent to the resistor value, R, and the extremum in ZIm occurs at a frequency = 2 f = 1/RC.

Ion transport can be modeled as a resistor or as a resistor in parallel with a capacitor (electrolyte geometric capacitance). Mass transport phenomena are fit using Warburg circuit elements, which represent diffusion processes. An example fuel cell equivalent circuit and Nyquist plots are shown in Figure 2.4.

In varying the bias from near open circuit conditions to very low cell potential, different processes can be seen to dominate the total cell impedance. If the test is repeated at multiple temperature, an activation energy for each identified process can be obtained, aiding in the identification of electrochemical or transport processes.

Figure 2.4 – Equivalent circuit model and Nyquist plots. (a) Example fuel cell equivalent circuit model and (b) corresponding Nyquist plots for different cell conditions (i, ii, and iii as indicated on inset I-V curve). As the current density of the cell increases, the reaction impedance loops shrink. It is not uncommon for one of the reaction loops to shrink by a greater percentage than the other(s). The Warburg element influence is not apparent until the highest current density curve (iii, green). The low-frequency intercept, high-frequency intercept, and the characteristic frequency for the cathode loop are shown for one of the Nyquist curves. The high-frequency intercept occurs at Re{Z} = R , the “ohmic” impedance due to charge transport. The low-frequency real axis intercept is the slope of the I-V curve.

While R||C is the idealized model for an electrochemical reaction, it is often necessary to use a constant phase element (CPE) in place of an ideal capacitor to improve the quality of the fit of the model to the data. The impedance of a CPE is Z( ) = 1/(CPE- T)(j )CPE-P, where CPE-T is in Farads, and CPE-P is a unitless constant. When CPE-P = 1, the CPE behaves exactly as an ideal capacitor. For many real fuel cells, the CPE- P value is less than 1. There are several suggested explanations for the physical origin of CPE behavior including compositional variation or porosity of the reaction surface, which may lead to a broadening of the range of reaction rates.

It can be very difficult to interpret EIS data without the support of additional experiments to isolate materials and morphological effects of one electrode. By comparing the test fuel cell results to a known control device, it becomes possible to associate design aspects with their respective impedance effects. Therefore, it is common to see comparisons of EIS data between several fuel cell assemblies in the literature.

2.6 Summary

Fuel cells are complicated electrochemical devices requiring both background knowledge of the physics and chemistry of their operation as well as careful analyses to fully understand the processes observed for any given device. This is true even for the relatively simple system of a H2/O2 fuel cell. More complicated reaction pathways can require a more intimate understanding of the chemistry and more thorough characterization methods than those discussed above, such as the inclusion of inlet and outlet stream gas analysis.

The techniques and parameters discussed above play an important role in understanding the results in the subsequent studies presented in this work. In particular, CV measurements were used extensively for the work presented in the bioelectricity chapter, while I-V and EIS measurements are discussed at length for the fuel cell experiments. All of these techniques may be used for both the bioelectricity device and for the more conventional fuel cells that are discussed; however, the

22 specific techniques used vary depending on the device of interest. In the bioelectricity studies, a three-electrode setup was used, which allows for easier analysis of CV results, while the fuel cell fabrication processes used by our lab do not easily facilitate a three-electrode setup. Instead, the fuel cell experiments consist of two electrode (no fixed-reference electrode), and potentials are given relative to the anode potential. Additionally, the I-V and EIS analyses would be more appropriate for the bioelectricity device in whole than in the in vitro studied on ferredoxin presented in §3. The reason similar techniques can be used for two different devices is that the behavior of each comes from the same thermodynamic principles. Differences in the thermodynamic or kinetic basis will be noted in the bioelectricity chapter.

3 Bioelectricity

This chapter will discuss a strategy for harvesting solar energy using photosynthetic organisms. The differences between this strategy and biological fuel cells will be discussed, as well as the challenges in implementation, and previous studies of ferredoxin, a protein essential to photosynthesis and bioelectricity. The simulations and experiments performed to study the anode reaction in the bioelectricity strategy are presented, as is a model for the efficiency of a photosynthetic bioelectric device.

3.1 Bioelectricity

The bioelectricity research performed by our group addresses the problem of solar energy conversion to electrical energy with a strategy to directly collect high energy electrons from the photosynthetic process. The schematic for this approach is presented in Figure 3.1. As shown in the figure, the biological system is interfaced with two electrodes at the scale of the photosynthetic system. One electrode inserted onto each side of the thylakoid membrane completes an electrochemical circuit with oxidation and reduction reactions separated by a membrane, creating an ion and voltage gradient between the anode and cathode. In this regard, the bioelectricity device is more closely related to a fuel cell than a photovoltaic device. To some extent, it may be considered a hybrid device with some characteristics of fuel cells and some of solar cells; sunlight is required to supply the energy input necessary for the cell to produce the “fuel”, ferredoxin, that is oxidized by the anode.

This section will begin by discussing the photosynthetic process, delineating some of the key differences between our proposed bioelectricity device and biological (microbial) fuel cells, and outlining the challenges of implementation and progress towards meeting these challenges.

Figure 3.1 – Bioelectricity schematic diagram and a simplified view of photosynthesis. Electrons energized by photosynthesis are transported by Fd to the anode inserted into the stroma of the thylakoid, passed through an external load, Z, and returned to the cell by the cathode reaction, resulting in regeneration of water in the thylakoid space. Typical thylakoid membrane thickness is approximately 5 nm. Figure not to scale.

3.1.1 Biological Fuel Cells

It is important both to recognize the field of biological, or “microbial”, fuel cells and to make clear the differences between such fuel cells and the bioelectricity studies carried out in our lab. Microbial fuel cells (MFCs) typically take advantage of anaerobic cellular respiration of organisms such as yeast and certain bacteria. An example arrangement, shown in Figure 3.2, places a current collector in solution with the organisms and a fuel in the anode compartment of the fuel cell. The fuel may be glucose or other biological compounds that can be broken down and used for energy by the organisms. In the absence of oxygen, the normal cellular respiration process (glycolysis) produces carbon dioxide, protons, and electrons. The protons diffuse through a semi-permeable membrane to the cathode compartment, and the electrons are shuttled to the anode electrode, typically by mediator chemicals also present in solution, although there are examples of microbial fuel cells that do not use such

25 mediators. The electrons are recombined with the protons in the cathode chamber, which is exposed to oxygen, to form water.

Microbial fuel cells are intriguing due to the abundance of possible fuels in the form of organic matter in wastewater and sediment. While this makes microbial fuel cells attractive, the power output is typically on the order of W/m2 or less. Large ion transport distances, including cellular membrane crossings, and mediating reactions lower the performance. There are MFC designs that reduce the ion transport distance and eliminate the need for mediators through organism selection. Even in these cases, MFCs are often process fuel in batches or with forced flow. In the latter case, the power input is often greater than that converted from fuel by the MFC, in which case the application is typically limited to wastewater treatment. Recent reviews of MFC technology are presented in the literature by Du (6) and Franks (7).

Figure 3.2 – Typical microbial fuel cell (MFC) schematic. Anode compartment may or may not contain mediator species that can cross cell membranes to transfer electrons to the anode. Protons diffuse through a membrane to the cathode compartment, which is oxygen-rich. The cathode compartment is sometimes open air.

The bioelectricity arrangement (Figure 3.1) eliminates the large ion transport distances by placing both the anode and cathode in the thylakoid of the cell, where extremely close proximity to Fd and the water splitting reaction would be maintained. By directly reacting with Fd and O2 in the chloroplast, the bioelectricity device can avoid mediator reactions, potentially yielding a higher efficiency from light absorption to current collection. We acknowledge that not all microbial fuel cells require light for operation, and that unlike traditional fuel cells, the bioelectricity device would likely benefit from pairing with a battery or other energy storage device (due to day/night cycles). It is possible, but not presently known, that a bioelectricity device may require replacement of algae support media or cells to maintain desirable current densities. We claim that by collecting electrons directly from photosynthesis, higher power densities than MFCs are achievable, and with sufficient advances in implementation, bioelectricity may become a viable alternative energy technology.

3.1.2 Photosynthesis

Plants, as well as algae and some bacteria, use photosynthesis to convert light energy into chemical energy that can be used by the organism to carry out the various cell functions.

Figure 3.3 shows the electron pathway in photosynthesis, with potential energy of the electrons indicated in the vertical axis. Electrons are obtained by splitting water molecules at a manganese complex in the thylakoid lumen. Each electron is initially excited by the energy gained from absorption of a photon in Photosystem II (PSII). The electrons are transported to Photosystem I (PSI) through a series of chemical steps in the thlyakoid membrane and excited by a second photon absorption event. At the end of the photosynthetic electron transfer chain is the mobile charge carrying protein ferredoxin. The ferredoxin (Fd) is reduced by the electron energized by PSI. At this point the reduced Fd has the highest reducing potential of any protein in the cell. The reduced Fd is the reductant for a larger protein, ferredoxin-NADP reductase (FNR), which in turn reduces NADP+ to NADPH, which is used in the Calvin cycle (3).

Figure 3.3 – Potential energy view of the photosynthetic electron pathway. The scale on the left indicates the electrical potential of the electrons relative to the normal hydrogen electrode. The thermodynamic voltage of bioelectricity is determined from the difference between the reduction potential of the water splitting reaction and the reduction potential of Fd. The low potential of water splitting (compared to standard potential of 1.23V) is due to the pH 5 environment in the thylakoid space.

The total change in potential energy of the electrons from the water splitting reaction to complexation with Fd is approximately 1.3 eV, slightly less than the energy from two photons absorbed by the chlorophyll due to the chemical reactions involved in the process. This energy difference is close to the thermodynamic cell potential of a hydrogen fuel cell. If one were able to effectively capture electrons from the mobile excited Fd in the chloroplasts, and return those electrons to the other side of the thylakoid membrane to react with protons and dissolved oxygen, a device similar in function to a fuel cell would be made (Figure 3.1). The major difference between this device and a fuel cell is that the “fuel” for the bioelectricity device is chemical energy converted from sunlight through photosynthesis, adding some steps to the overall system process. However, the basic reaction thermodynamics and kinetics that determine the behavior at the two electrodes are essentially the same as those

28 discussed in §2. Additionally, the mass transport behaviors are similar, the principle difference being liquid solution versus gas phase. Ion transport in the bioelectricity device and solid oxide fuel cell electrolytes can be treated the same (i.e. use same equations), but have significantly different mechanisms; liquid electrolyte diffusion versus solid electrolyte vacancy “hopping”. Rather than “fuel utilization”, , it may be more pertinent to consider the efficiency of the organism‟s ability to convert sunlight to charged Fd along with the fraction of charged Fd proteins oxidized by the inserted electrode and not used for cell processes. In this case, the net device efficiency may be given by

 Grxn Vcell N Fd net Atot fanode 3.1 H HHV Ethermo abs  where Atot is the total fraction of light absorption by the device, N Fd is the number of

Fd molecules charged by PS I per unit time, Fd is the number of photons absorbed per unit time, and fanode is the fraction of Fd oxidized by the anode (and not by FNR). Because two photons are required to produced one charged Fd, the net efficiency is therefore necessarily less than 50%.

It is interesting to note that the pH of the thylakoid stroma and lumen plays a role in determining the theoretical cell potential of the bioelectricity device. The pH of 5 in the lumen lowers the theoretical potential of the water reaction. The reaction potential for Fd oxidation is approximately –420mV vs NHE at the pH of 8 in the stroma (8). This pH gradient is maintained by active pumping by proteins embedded in the thylakoid membrane. It is not clear how the bioelectricity scheme might affect the proton pumping or other cell functions. We suspect that the diversion of reduced Fd from the FNR complex to the inserted anode will serve to limit the cell function by lowering the efficiency of conversion from light energy to chemical energy useable by the cell.

3.1.3 Challenges

There are significant challenges that must be met in order to implement the bioelectricity scheme shown in Figure 3.1. Among these challenges are nanoprobe design and fabrication, organism immobilization, and insertion of the probes into the cell.

The thickness of the thylakoid membrane is on the order of 5nm. This length scale limits the size of the nanoprobes at the tip and, to some extent, the spacing between the anode and cathode. The probes must be electrically conductive at the tips so the anode and cathode reactions can occur, and electrically isolated from each other to ensure that the electrons obtained by the anode flow through an external circuit and not an electrochemical short to the cathode. Furthermore, the probes must be able to penetrate the cell membrane, chloroplast envelope membrane, and thylakoid membrane. To meet these demands, nanoscale probes were fabricated from silicon wafers using a combination of thin film and photolithographic process. The probe fabrication and testing processes are described by Bai in the literature (9). The probes consist of a silicon nitride cantilever surrounding a thin film of conductive metal (Au). The outer surface of silicon nitride lends mechanical strength as well as electrical insulation. The cantilever is cut to a high aspect ratio and split into the two separate probes with focused ion beam milling. The tips of the probes are milled to expose the underlying gold layer, resulting in a twin-tip ultra sharp probe with electrical access to the tips.

For the initial experiments we elected to use the alga Chlamydomonas reinhardtii. This organism was selected for its relatively large size (~5-10 m diameter), large chloroplast size relative to the cell body, and the availability of a wide array of mutant varieties. C. reinhardtii, shown schematically in Figure 3.4, uses a pair of flagella for cell motion. To insert nanoprobes into a live cell, the cell must be immobilized. Ryu et al demonstrated immobilization by flowing media containing C. reinhardtii mutants without flagella through a microfluidic channel with triangular “traps” that allow the solution to flow but catches the cells (10). A confocal microscope was modified so

30 that an atomic force microscope (AFM) could be used to control the nanoprobe x-y positioning and distance from the cells in the z-direction. The AFM is advantageous as it potentially allows for more accurate placement of the probe tip during insertion into the cell through scans of the surface to identify the chloroplast location. Additionally, the force feedback is beneficial during insertion as it helps to eliminate guesswork of reaching an appropriate depth in the cell.

Figure 3.4 – Schematic drawing of Chlamydomonas reinhardtii. Note the size of the chloroplast relative to the cell body, as well as the flagella used for cell motion. Nanoprobes are shown inserted into the thylakoid.

Once inserted into the cell, it is important to determine if a photo-induced current can be captured by the probe. Indeed, Ryu and other members of our lab were able to demonstrate an oxidation reaction current at the probe tip inside a cell, and furthermore showed that the current was due to photosynthetic activity, although it was not clear what reactions may have been occurring at the electrode (11).

These results, while promising, do not address the question of Fd oxidation; it is not clear if the current was due to oxidation of Fd or another species. In fact, it is generally accepted that Fd oxidation exhibits poor kinetics on a bare metal electrode (12). This is found to be true even for surfaces such as platinum which are often catalytically

31 active. The next few sections of this chapter will deal with the chemistry of Fd in the cell and Fd oxidation by metal electrodes.

3.2 Ferredoxin

The metalloprotein ferredoxin is a small (~11.6 kDa) iron-sulfur protein that serves as an electron carrier for a number of biological redox reactions. In photosynthetic organisms, it mediates the reduction of NADP+ by coupling electron transfer from photosystem I to the ferredoxin-NADP reductase (FNR) enzyme (13). Two ferredoxin (Fd) molecules are required to fully reduce the FNR prosthetic group, flavin adenine dinucleotide (FAD). The structure of Fd and FNR have both previously been determined [PDB citations], and reduced Fd carries its extra charge in the [Fe-S] core. In the case of C. reinhardtii, the core contains two Fe atoms and two S atoms. The Fe atoms are also bonded to the S moiety of cysteine groups, and are thus each tetrahedrally coordinated to four S atoms. The structure of the Fd-FNR complex is shown in Figure 3.5a, and the oxidized and reduced forms of FAD are shown in Figure 3.5b.

The redox chemistry of Fd has been studied by many groups, and it has been found that the reduction of Fd is irreversible at bare metal and carbon electrodes (12; 14). It should be noted that there have been reports of reversible oxidation of Fd with the use of mediators in solution (14; 15; 16; 17) as well as modified carbon (18; 19) and In2O3 electrodes (20). Based on these prior reports and taking inspiration from what is known about Fd redox in nature, a metal electrode coated with a FAD analog may be able to improve the kinetics of charge transfer relative to the untreated electrode. The first structure considered for this was 4-mercaptopyridine (4-PyS), which has a simple ring structure with a N site, and a S moiety, which enables self-assembly of the molecule on a Au(111) surface. Initially we neglected to consider the pKa of the surface modifier and the requirement that the N sites in FAD be protonated to NH+ before they can be reduced. This oversight resulted in some apparent discrepancy between simulations on 4-PyS and experimental results.

Figure 3.5 – Fd-FNR complex and FAD redox. (a) Zea mays Fd docking with FNR (PDB 1GAQ,(21)), with [2Fe-2S] and FAD visible, and (b) FAD reduction to FADH2.

3.3 Simulations

In tandem with the experimental work exploring surface functionalization, quantum chemical and molecular dynamics calculations were performed to predict the relative efficacy of charge transfer between Fd and Au(111), H+PyS-Au(111) and FADH+. Quantum chemical simulations were used to predict the electron affinity and electronic coupling matrix elements for the three systems. Molecular dynamics simulations were performed to study possible orientation effects of positively charged (i.e., protonated) surfaces on the negatively charged Fd. The specifics of the methods used are detailed + in (22). The Au(111) surface was simulated as a Au10 cluster. The H PyS-Au(111) + used Au10 for the electrode surface. FADH was approximated by isoalloxazine with a

CH3 functional group (R in Figure 3.5b) in place of the remainder of the riboflavin and the nucleotide of the full FAD molecule. This section presents a summary of the key results and their implications towards the experiments. Figure 3.6 shows the structures used for the quantum chemical calculations.

Figure 3.6 – Structures used for quantum chemical simulations. + + Top left, Au10 with H PyS; Bottom left, isoalloxazine used for FADH ; Right, [2Fe-2S] core region of Fd and H+Py. Note that the „R‟ group of FAD (see Figure 3.4) is replaced with a methyl group in this image. Nitrogen sites marked „A‟ and „B‟ are the first and second reduction sites, respectively. S (yellow), N (blue), C (black), O (red), H (grey), Au10 (shaded blue in top left)

Density functional theory calculations showed similar electron affinities (~5eV) for FADH+, and H+PyS, while the Au cluster electron affinity was significantly lower (~2.3eV). The ionization potential of Fd is approximately 5eV, therefore reduction of FADH+ or H+PyS to FADH or HPyS, respectively, would be able to provide sufficient energy to liberate the electron coordinated with the [2Fe-2S] core of Fd. This suggests that H+PyS is a reasonable candidate for electrode functionalization to oxidize Fd.

Configurational interaction calculations were used to determine the relative probability of electron transfer, and therefore a qualitative indicator of reaction kinetics, between reduced Fd and Au(111), H+PyS-Au(111), and FADH+. The literature (23) suggests 2 that the rate of electron transfer scales with Hif , where Hif is the off-diagonal Hamiltonian matrix element representing the interaction of Fd and the acceptor (gold,

34 mercaptopyridine, or FAD), with an electron localized on either Fd or the acceptor. Using a single excitation configurational interaction method (CIS), it was found that the coupling matrix element Hif was 0.2 eV, 0.4 eV, and 0.75 eV for a bare Au surface, H+PyS, and FADH+, respectively. Normalizing to the bare gold result as a benchmark, the relative electron transfer probabilities as determined by comparing 2 + + values of Hif are 1, 4.0, and 14.1, for Au, H PyS, and FADH , respectively. This result suggests that the H+PyS is a significantly better electron acceptor for reduced Fd than a bare gold surface, although it is still significantly worse than the system that occurs in nature.

It should be noted that the method for determining electron transfer probability assumes the weak coupling limit is valid. However, the Hif values are quite large for the assumption to hold well. In the strong coupling limit, electron transfer is assured if the electron affinity of the acceptor is equal to or greater than the ionization potential of the donor. If the acceptor electron affinity is less than the donor ionization potential, the difference in energy serves as an activation energy barrier that can be lowered by geometry changes. The simulations did not exhibit significant geometry change with electron transfer, suggesting that other considerations be more important in determining electron transfer probabilities. (22)

Molecular dynamics simulations showed that the NH+ residue of a pyridine ring appears to help orient the charged Fd so that the [2Fe-2S] core is closer to the electrode surface. This is qualitatively consistent with the orientation of Fd when it is complexed with FNR.

3.4 Experiments

Improving electron transfer to an electrode for the bioelectricity scheme is the end goal of this work. This section presents the results of experiments performed in an attempt to improve heterogeneous charge transfer from Fd to a conductive electrode. A clean gold surface is the reference point for benchmarking the efficacy of different electrode surfaces. The results of this work has been published in the Journal of the Electrochemical Society (24).

3.4.1 Materials and Methods

Ferredoxin from Spinacia oleracea (spinach, [2Fe-2S]), 2.2 mg/mL in 0.15M Tris with NaCl, pH 7.5; 4-mercaptopyridine (PyS), 95%; cystamine dihydrochloride (2- aminoethanethiol dihydrochloride, AET), 98%; poly-l-lysine hydrobromide with average molecular masses of 15 kDa and 24 kDa (PLL15 and PLL24, respectively); 11- mercaptoundecanoic acid (MUA), 95%; triethanolamine hydrochloride (TEA), 99+%; 1-ethyl-3-[3-(dimethylamino)propyl]carbodiimide hydrochloride (EDC); N- hydroxysulfosuccinimide (NHSS), 98.5%; and sodium dithionite, 85%; were purchased from Sigma-Aldrich and used without further purification. Platinum film, 99.9%, was obtained from Sigma-Aldrich and cut into strips for use as counter electrodes. A AgCl/Ag reference electrode and 3 M KCl were purchased from Microelectrodes, Incorporated (Bedford, NH, USA). All electrochemical experiments were performed in an anoxic N2 environment (< 1 ppm O2) containing 1-2% H2. Planar Au electrodes were fabricated by evaporation to form a layer 200nm thick on silicon wafers. Before Au deposition, an adhesion layer of approximately 5nm of titanium was deposited by evaporation. The ITO was a gift from the laboratory of Michael McGeehee. X-ray photoemission spectroscopy (XPS) revealed that the ITO was approximately 3 at% Sn. The geometric area of all electrodes was 0.32 cm2.

Fd was used, undiluted, as received from Sigma-Aldrich. The concentration of 2.2 mg/mL corresponds to approximately 190 M, based on a molecular weight of 11.6 kDa. We did not observe any variation in results between the Fd solutions we received. The dithionite, which serves as a reductant, was dissolved in 0.15 M Tris-

HCl, pH 7.5 with 0.25 M NaCl. The dithionite and Fd were mixed in a 1:1 molar ratio, and the excess dithionite was removed from the reduced Fd by repeated centrifugation and dilution using a NanoSep 3K filter manufactured by Pall Corporation (East Hills, NY, USA). To verify the efficacy of the Fd reduction process, a sample containing the reduced Fd solution was then sealed in a cuvette for photospectrometry. After a wavescan on the reduced solution, the cuvette was opened and exposed to air. The Fd was then oxidized by blowing air onto the surface of the solution. A second wavescan revealed an absorption peak at 420nm that had not been present in the reduced solution. The appearance of this peak after oxidation is consistent with reports in the literature (25).

3.4.2 Electrode Surface Modification

The species used for surface functionalization are shown in Figure 3.7. Initially, small molecules with amine and thiol moieties were chosen to mimic the NH+ reaction site of FADH+ and to allow for self-assembly of ordered monolayers on a gold surface through the strong Au-S bond. Self-assembled monolayers (SAMs) have been studied extensively and numerous reports can be found in the literature. The underlying mechanism tends to be a highly energetically favorable covalent bonding of molecules to a surface, such as alkanethiols on gold or silver. The (111) surface of Au or Ag offers low-energy bonding sites in the nexus between three surface atoms (eg, R-S-

Au3). The ordering of the SAM occurs to close-pack the molecules bonded to the surface and the arrangement is dependent on deposition conditions and interactions between molecules (26). Later experiments investigated adsorption of polymeric L- lysine on an ITO electrode, as well as a more involved process for Au functionalization with poly-L-lysine.

Figure 3.7– Structures and pKa values of surface coatings. The above molecules were used to functionalize the electrode surfaces, with relevant pKa values indicated. Shown in unprotonated state, from left to right, are 4-mercaptopyridine, 2-aminoethanethiol, and poly-L-lysine.

3.4.2.1 Mercaptopyridine

4-mercaptopyridine was chosen as a first approximation to the redox active region (site „A‟ and surroundings in Figure 3.6) of FAD. The quantum chemical calculations predicted that the electronegativity of H+PyS is comparable to that of FAD and that H+PyS on gold is expected to exhibit a 4-fold increase in oxidation efficiency, assuming the results of the configuration interaction simulations are valid.

PyS-Au electrodes were prepared by immersing a Au electrode into ethanolic, aqueous, or 0.05 M perchloric acid solutions of 1 mM PyS for 1 minute followed by rinsing of the electrode by immersion in ethanol, water, or 0.05 M perchloric acid, respectively; the rinsing step, which removes the excess PyS that adheres to the surface, always used the same solvent as that used for SAM formation. This process is described in the literature (27).

XPS analysis of the PyS-Au electrode used a model with a Au underlayer and a surface layer with the composition of the PyS molecules (C5NS). The photoemission spectrum of the sample was measured at both 30° and 90° off-plane, near the primary

38 binding energies for Au, C, N, S, and O. Typical data showed the large Au underlayer and the presence of N and S atoms on samples that had received the SAM treatment. The peak for N 1s emission was centered at approximately 401.3 eV, and the peak for

S 2p3/2 emission was centered at 163.2 eV, which is consistent, within 0.5 eV, with the energy associated with S bound to a gold surface (28). The N emission peak is in agreement with literature values for protonated pyridine (29). Samples that were not treated with the SAM solution did not show any measurable emission at the energies associated with sulfur. The thickness of the C5NS layer was calculated from the data to be about 2 Å, which is consistent with the length of the PyS molecule and the tilt observed in packing (30; 31; 32), indicating that a monolayer of PyS had formed on the Au surface. The contact angle for water on a 200 nm thick Au surface before SAM treatment was found to be 90±1°. The freshly applied PyS monolayer reduced the contact angle to 20±1°.

At pH 7.5, the PyS-Au electrodes yielded CV results similar to those observed with a bare Au electrode. Quantum chemical calculations (§3.3) for a protonated pyridine ring on Au10 predicted that oxidation of a [2Fe-2S] cluster, such as the cluster associated with reduced Fd, should be approximately 4 times more favorable on H+PyS-Au than on bare Au (assuming the weak coupling limit assumption holds).

Protonation of 4-PyS is expected at a pH lower than its pKa, which is approximately 5.5 (33). We initially hypothesized that the absence of the 4-fold increase in the rate of electron flow from Fd to the 4-PyS surface of the electrode was due to the lack of protonation of the 4-PyS in the pH 7.5 solution. Based on the pKa, only about 1% of the PyS is expected to be protonated, as determined by the Henderson-Hasselbalch equation:

[A - ] pH pK log a [HA ] 3.2 When the working solution is titrated with HCl to pH 4.5, we can expect roughly 90% protonation of PyS to H+PyS, and we observed stronger oxidation and reduction peaks associated with Fd, with the anodic peak at a potential of approximately –500 mV and

39 a peak-to-peak separation of 80 mV for both Au and H+PyS-Au. The reduction peak potential is more positive than expected (8), suggesting that the proton concentration plays a role in Fd electrochemistry at pH values close to the isoelectric point of Fd, which is approximately 4 (34). Figure 3.8 shows typical results for PyS-Au and Au electrodes in a 2.2 mg/mL Fd solution. CV curves were also measured at v = 10 mV/s, for which the peak current scaled as approximately v0.8 in the acidic solution, and scaled with v0.5 at neutral pH, suggesting that, despite appearances, the neutral case is closer to diffusion-control than the acidic case.

3.4.2.2 Aminoethanethiol

AET-Au electrodes were prepared by immersing a Au electrode into an aqueous solution of 0.02 M cystamine dihydrochloride for 2 hours, followed by rinsing of the electrode in water. This process has been previously performed as part of a study of pyrroloquinolinequinone deposition on Au (35).

AET was chosen as the second SAM molecule to be tested with a Au electrode. As previously mentioned, oxidation of Fd has been observed with an electrode consisting of positively charged alkane molecules adsorbed to a carbon surface. AET is a short- + chain alkanethiol that has a NH3 moiety in neutral solution (36). We expect that short alkane chains may be more desirable than larger molecules for the penetration of chloroplasts by nanoprobes because of the lower chance of mechanical damage to the cell and the SAM during penetration. XPS analysis of the AET-Au electrodes showed

N 1s emission at 400.6 eV, and S 2p3/2 emission at 162.1 eV, which is consistent with + published results for CnS bound to Au (37). Although the NH3 moiety should encourage Fd-associated redox activity, none was observed in the CVs using the AET- Au electrodes, suggesting that the positively charged surface layer is not a sufficient condition for Fd oxidation to occur.

Figure 3.8 – PyS-Au results. Cyclic voltammograms for: Fd at pH 7.5 for Au (a, black) and PyS-Au (b, blue) electrodes; and Fd at 2 pH 4.5 for Au (c, gray) and PyS-Au (d, red) electrodes. Scan rate 20mV/s. Electrode area 0.32 cm . Positive current is reducing.

3.4.2.3 Poly-L-lysine

Au electrodes were coated with PLL by amide bond linkage to a MUA SAM. Au electrodes were immersed in MUA for at least 24 hours, and then placed in an aqueous solution containing 70 mM EDC and 12 mM NHSS for 1 hour. The electrodes were then placed in pH 8.0 TEA buffer containing 1 mg/mL PLL of molecular mass of either 15 kDa or 24 kDa for approximately 90 minutes. The PLL-Au electrodes were rinsed in 0.1 M NaOH for 20 minutes. This process was described previously for adsorptive biosensor applications (38).

The contact angle of water with the Au electrodes treated with MUA was 34±4º, and increased to 54±1º after attachment of PLL to the gold electrodes. The PLL15-Au and

41 the PLL24-Au electrodes both facilitated reversible oxidation of Fd. Typical CV results from the PLL-Au electrodes are shown in Figure 3.9. The CV for PLL15-Au shows a second redox pair observed for the PLL-coated electrodes. The peak current in this smaller pair was less than 500 nA and was observed using both the PLL-Au and PLL- ITO electrodes, suggesting that the redox pair is associated with PLL. CV curves were also measured at v = 10 mV/s, for which the peak current scaled as approximately v0.5. The difference in Fd cathodic peak current may be due to differences in the extent of initial reduction of Fd by dithionite.

Figure 3.9 – PLL-Au results. PLL15-Au (a, blue) and PLL24-Au (b, red) electrodes in reduced Fd. Scan rate 20mV/s. Electrode area 0.32 cm2. Positive current is reducing.

Figure 3.10 – PLL-ITO results. Bare ITO (a, grey), PLL15-ITO (b, blue), and PLL24-ITO (c, red) electrodes in reduced Fd. Scan rate 20 mV/s. Electrode area 0.32 cm2. Positive current is reducing.

The contact angle of water on untreated ITO was measured as 84±1º. After a 10- minute immersion of the ITO in either PLL15 or PLL24 in 0.15 M Tris-HCl, pH 7.5, the contact angle dropped to 69±2º for ITO electrodes coated with either PLL15 or PLL24, suggesting a different arrangement of PLL on the surface than in the case of the PLL- Au electrode. XPS was performed on a sample of untreated ITO to determine the level of doping. Analysis of the relative areas of In, Sn, and O peaks indicated that the sample was approximately 3% Sn doped

In2O3. Both of the PLL-coated ITO electrodes facilitated reversible oxidation- reduction of Fd, but the reactions were weaker than those observed using the PLL15- Au electrodes. The charge transfer coefficient and, when normalized for Fd concentration and electrode area, peak currents for the PLL-ITO electrodes were similar to that of the PLL-In2O3 electrodes, as previously reported (20), although our results show a larger peak separation. Typical results for PLL-ITO electrodes are

43 shown in Figure 3.10. CV curves were also measured at v = 10 mV/s, for which the peak current scaled as approximately v0.5.

3.4.3 Analysis of results

Table 3.1 summarizes the results of the CV experiments and reports the transfer coefficient, , and the standard heterogeneous rate constant, k0. We assume that the published value for the diffusivity of Fd in water of 11.9±0.7x10-7 cm2/s is valid for Fd in 0.15 M Tris (39). The rate constant was calculated using the method of Nicholson

(40), as a function of the separation between the anodic and cathodic peaks, ( Ep), that follows the empirical relation

0 F DR k ( E p ) DO v 3.3 RT DO where DO is the diffusivity of oxidized Fd, DR is the diffusivity of reduced Fd, and v is the scan speed. We also assume that the diffusivity of Fd is the same for the reduced and oxidized states. This eliminates the transfer coefficient from Equation 3.3. Although Nicholson assumes that = 0.5, slight asymmetry in the free energy barriers to charge transfer should not affect the value of , since Ep is approximately independent of for 0.3 < < 0.7. To verify that the transfer coefficient is in this range, the data were fit to a linear regression of the logarithm of the current using the expressions

lni lni0 F / RT (O e R) 3.4 lni lni0 (1 ) F / RT (R O e)

0‟ 0 where = E – E , and i0 is the exchange current, which can be related to k by

0 1 i0 FAk CO CR 3.5 * * where CO and CR are the bulk concentration of the oxidized species (O) and reduced species (R), respectively. The expressions in Equation 3.4 are the Butler-Volmer equation (Equation 2.12) for cases where charge transfer is either reductive or oxidative, but not both, rewritten in log form. This approximation is valid when the

44 surface concentration of one species is small relative to that of the other. Therefore, when fitting the CV data, we only considered the portion of the curve at potentials well before the inflection in current due to diffusion limitations, which appears linear in the kinetic limitation regime. We calculated the transfer coefficients using the slope of the natural logarithm of the current rise at 15% of the peak current. Note that the Au and PyS-Au electrodes were found to have transfer coefficients outside the range for which the Nicholson method for estimating k0 is independent of . Therefore, the rate constants we report for these electrodes may overestimate the true rate constant. Additionally, the oxidative waves of the CV curves for Au and PyS-Au at pH 7.5 appear to have an adsorption controlled shape. However, the estimated value of

( Ep) and the curve shape are in agreement with that for slow (nearly irreversible) quasireversible diffusion-controlled charge transfer (5; 41). It is necessary to note that, in addition to the transfer coefficient falling outside the range of 0.3< <0.7, there is an uncertainty in the estimated k0 for these electrodes due primarily to the extrapolation of Nicholson‟s empirical results to obtain ( Ep) for the peak separation of 250 mV. The value of this uncertainty exceeds the magnitude of the rate constant, suggesting that this method of analysis is not valid for large peak separations.

Table 3.1 – Summary of CV results and analysis.

4 0 ipa / Epa / mV 10 k 1 – Electrode E / mV -1 A vs AgCl pa / cm-s (reductive) (oxidative) Au, pH 7.5a 1.3 -450 250 1 ± 3b 0.3 0.2 PyS-Au, pH 7.5a 1.2 -470 245 1 ± 3b 0.3 0.2 Au, pH 4.5 2.6 -500 80 20 ± 2 0.8 0.2 PyS-Au, pH 4.5 2.9 -500 80 20 ± 2 0.7 0.3 AET-Au ------PLL15-Au 2.4 -605 105 9 ± 2 0.6 0.5 PLL24-Au 2.7 -615 110 8 ± 2 0.6 0.4 ITO ------PLL15-ITO 1.7 -590 160 3 ± 0.5 0.5 0.5 PLL24-ITO 1.2 -590 155 3 ± 0.5 0.6 0.5 All data for scan rate of 20 mV/s. a May not be diffusion-controlled charge transfer. b Estimates based on extrapolation of Nicholson‟s data. Large peak separation makes these values unreliable and likely overestimates of the true rate constant.

The values for k0 determined from Equation 3.3 are similar to those obtained by regression and Equation 3.5. In estimating the rate constant with Equation 3.5, we used the geometric area of the electrode. We report the values of k0 as determined by Equation 3.3. The values for were interpolated from the data generated by Nicholson (reproduced in Table 6.5.2 of (5)). Peak currents are reported after subtracting background curves measured in Tris buffer without Fd, and applying the following method, (5), where ipc is the peak cathodic current, isp0 is the current at the switching potential, and ipa0 is the anodic current relative to the cathodic baseline, to determine the true anodic current, ipa:

ipa ipa0 0.485isp0 0.086ipc 3.6 We found the peak anodic currents to be accurate to within about 0.1 A. The peak-to- peak separation, Ep, is given to the nearest 5 mV; a resolution of 5 mV was used when measuring the curves.

Because of the uncertainty in the values we estimated for the transfer coefficient for the reduction and oxidation processes, the values for and 1 – do not always add to unity. In particular, the formal redox potential for Fd is very close to the H2 formation limit of the electrode window, and thus in order to accurately calculate 1- from the CV data, a baseline must be estimated from not only the reduction current, but also the observed H2 formation current. Subtracting the background from a scan on 0.15 M Tris-HCl, pH 7.5 helps determine this baseline, but is not completely accurate due to the decay of the reduction current due to diffusion and variations in rates of H2 formation. Because the baseline for the reductive wave is straight-forward and that for the oxidative wave is more difficult to obtain, the reductive transfer coefficients are likely more accurate than the oxidative. One should always assume that uncertainty is associated whenever a transfer coefficient is reported due to the difficulty in extracting this parameter from real data without simulation, for which there are many parameters that must be accurately fit in order to obtain a reliable value for . In the case of Au and PyS-Au at neutral pH, the sum of and 1 – is significantly less than unity. This suggests that the Butler-Volmer model for the current rise near E0‟ used to determine

46 the charge transfer coefficient is not valid for the Au and PyS-Au electrodes at pH 7.5. Qualitatively, this low sum and the rate constant imply that both bare Au and PyS-Au electrodes do not reversibly oxidize reduced Fd at pH 7.5.

At pH 7.5 there was a negligible change in the kinetic rate constant for electron transfer from Fd to the PyS-Au electrode relative to the bare Au electrode; the addition of the PyS monolayer did result in an increase in at pH 4.5. The rate constants found for the Au and PyS-Au electrodes at pH 4.5 suggest that the role of the solution pH in Fd redox is more significant than protonation of the PyS monolayer. Furthermore, despite the finding that a neutral pH favors the protonation of the AET-Au electrode, no Fd redox is observed. This suggests that although simulations show that protonated pyridine increases the likelihood of charge transfer from reduced Fd, and positively charge dodecylamine on a carbon electrode has been shown previously to oxidize Fd (18), a protonated nitrogen moiety is not a sufficient condition for charge transfer to + occur. Lysine has a pKa of ~10.2, indicating that PLL is likely completely in the NH3 state at neutral pH (3). Additionally, the shape of the PLL surface may allow for closer + proximity between NH3 of PLL and the [Fe-S] core of Fd.

3.5 Bioelectricity Device Calculations

This section presents a simple light absorption model for C. reinhardtii used to estimate the fuel cell behavior of a hypothetical bioelectricity device comprised of a 1 m2 array of cells. The anode kinetic losses for such a device are estimated using the kinetic parameters from the S. oleracea Fd experiments on Au and PLL-Au electrodes.

3.5.1 C. reinhardtii photoabsorption model

Using spectrophotometric transmission data for a solution of C. reinhardtii, we use the density of cells in solution to approximate a pseudo-close packed area density and determine how many pseudo-close packed planes with this density of cells the light passes through in a 1 cm cuvette. Figure 3.11 illustrates the geometry assumed for the light absorption model. To start, we define an effective radius, reff, of solution space occupied per algal cell based on the inverse of cell density in solution. Assuming

47 close-packed spherical volume per cell and ignoring the error from the space between spheres, the effective radius is

1/ 3 3 r 3.7 eff 4 where is the number of cells per unit volume, approximately 2x105 cm-3 in the sample used for absorption measurements. The distance between nearest-neighbor algal cells is 2∙reff.

Figure 3.11 – Photoabsorption model geometry. Model assumes normal incidence of light for close-packed spheres of solution with radius reff for each cell (green circles). The spacing between close-packed planes, d, is also shown. In the image above, monochromated light enters the solution from the left, and total transmission is measured on the right by the photodetector.

Next we need to determine the cross-sectional area density of cells, eff, per close- packed plane. The quantity eff represents the ratio of the cross-sectional area of cells with radius rcell to the area of a unit cell of close-packed spheres in a plane; this is the fraction of the plane area that absorbs light and is expressed as a function of the cell radius and reff:

2 rcell eff 3.8 2 3 reff The distance between planes, d, can be written as

2 6 d r 3.9 3 eff Lastly, the number of planes of cells that light passes through in the cuvette is given by (1 cm)/d.

Because we are interested in the absorption of a single layer of cells in an array, we must back out the absorption of one plane of cells in the solution from the total absorption. The absorption of one layer can be given by

1/ N A1 1 (1 Atotal ) 3.10 where Atotal is the total absorption by the solution in the cuvette. This equation comes N from the fact that the total transmission, Ttotal, is equal to T1 and A = 1 – T, where T1 is the transmission of a single layer of cells. Notice that this neglects scattering effects. We now need to be able to scale absorption up so that it is consistent with the fractional absorber area of the scheme put forth in this proposal. If we build an array that fixes cells with circular cross-sections in a square array shoulder-to-shoulder, then /4 (or about 78%) of the array area is absorber material, regardless of the size organism we choose. The scale factor for absorption of our proposed array is then

2 3 reff 3.11 4 eff 2 rcell Now we can write out the absorption for a stack of n arrays of cells, using the scale factor above to estimate the absorption of our proposed array from spectrophotometry data:

2 n 3 reff An,array 1 1 A1 3.12 2 rcell

Figure 3.12 shows absorption as a function of wavelength for stacks consisting of n = 1 through n = 5 arrays of cells. A stack with 5 absorbing layers yields 90% total absorption of the solar spectrum (42), or about 1932 mol photons m-2 s-1.

Figure 3.12 – Estimated bioelectricity device photoabsorption. Absorption of stacks of square arrays of C. reinhardtii estimated from spectrophotometric measurements of algal solution. Number of layers in stack indicated by color: 1, blue; 2, red; 3, green; 4, cyan; 5, magenta. Black curve is solar spectrum (42). A stack of 5 layers results in roughly 90% absorption of sunlight.

3.5.2 Maximum current density

The solar constant is approximately 1.4 kW/m2 irradiance. However, this considers the total energy of each photon and all wavelengths of light, most of which are outside the spectrum utilized by the photosynthetic reaction centers. A more accurate model for light available to photosynthesis considers only the number of photons on the range of 350 nm to 750 nm light. This value is approximately 2142 mol photons/m2-s. If the cells absorb every photon, we would see a maximum achievable power of about 140 W/m2. This estimate uses the thermodynamic potential of 1.35V discussed earlier and the fact that 2 photons are needed to excite 1 electron from water to reduce ferredoxin. Taking into account absorption data for C. reinhardtii and the model discussed in §3.5.1, we find that a stack of 5 arrays of cells would absorb 90% of the incoming light, and would thus have a maximum power output of 126 W/m2. It is

50 important to recall that 1.35V is an open circuit voltage (OCV) and cannot be achieved when there is a nonzero current flow. A more realistic value for maximum power output can be obtained by using an operating potential of approximately one-half the OCV; traditional fuel cells typically operate in this regime. Therefore, we can reasonably expect to obtain an operating potential of 0.7V and a maximum power output of 65 W/m2. By incorporating information obtained through the analysis of experimental results in §3.4, a more accurate estimate may be made.

3.5.3 Bioelectricity device fuel cell behavior

Here we expand the bioelectricity device model to include anode kinetic losses, which can be estimated using the kinetic information obtained in the S. oleracea Fd experiments. From the heterogeneous rate constants, k0, determined by the experiments discussed in §3.4, one can estimate the exchange current density for the anode reaction of the hypothetical bioelectricity device by

0 j0 Fk C 3.13 where F is Faraday‟s constant and C is the concentration of reduced Fd in the cell environment. Taking the Tafel approximation, the overpotential required to maintain a given current density can be determined as in Equation 2.17 using the value for the transfer coefficient, , as determined in §3.4. The anode activation loss is shown in Figure 3.13 for both bare gold and PLL-Au electrodes. In the figure, we have assumed a Fd diffusion-limited current equal to the maximum theoretical current from photoabsorption. This requires assumptions regarding the diffusion length of Fd to the electrode. Using Equation 2.22 for limiting current density,

o eff cFd jL nFD 3.14 we can use the diffusivity of Fd from the literature for Deff and an estimate for concentration of reduced Fd in the cell to yield an estimate of the diffusion length, , on the order of 10 nm. This assumes that approximately 50% of the Fd in the cell is in the reduced state.

Figure 3.13 does not include the effects of any ohmic losses or cathode activation losses from the water formation reaction. The ohmic and cathode losses will necessarily reduce the output from the estimate, but the important point is the marked increase seen in fuel cell performance for the hypothetical bioelectric fuel cell with the PLL coating on the anode.

Figure 3.13 – Projected fuel cell behavior for hypothetical bioelectricity device. Based solely on anode activation losses. OCV determined from reduction potentials of anode and cathode reactions. Limiting current determined from 90% light absorption. Charge transport and cathode kinetic losses not considered.

3.6 Conclusions

We have studied the oxidation of reduced spinach Fd using a three-electrode CV arrangement and a variety of working electrodes with the goal of increasing charge transfer to the anode of the bioelectricity from photosynthetic electron flow scheme outlined in Figure 3.1. We found that both gold and ITO electrodes that were modified by covalent attachment and adsorption, respectively, of PLL of 15 kDa and 24 kDa average molecular mass served to improve the kinetics of Fd oxidation. Based on the results of this experiment and the relative ease of nanoprobe fabrication utilizing Au as

52 compared to ITO, the PLL-Au electrode offers the most promise for facilitating in vivo oxidation of Fd.

However, calculations of photoabsorption and anode activation losses result in an estimate for the maximum possible power density that, while of practical use, does not compete with commercially available solar technologies. This is a significant detriment to the development of a bioelectricity device, particularly since the development of such a technology as a practical device is likely decades away. It may be possible to increase the attractiveness of a bioelectricity device through alternate approaches to cell-at-a-time penetration. For example, it may be possible to coax chemically opened cells (ie, with lysed membranes) into attaching to a substrate with properly designed surface coatings. In some primitive organisms, such as Gloeobacters, the photosynthetic structure resides in the cell membrane (43), and reduced Fd is mobile in the matrix of the cell body, and not restricted to the chloroplast stroma as in C. reinhardtii. In this arrangement, the organism would seal around the electrode, making the substrate fabrication relatively simple (and less expensive).

4 Low Temperature Operation of YSZ DMFC

This chapter deals with a direct methanol fuel cell (DMFC) with an yttria-stabilized zirconia (YSZ) electrolyte, an oxide ion conducting ceramic. Typically, DMFCs use a polymeric electrolyte and operate at temperatures around 100°C. By using a ceramic electrolyte, some of the problems of polymer electrolytes are eliminated. However, ceramic electrolyte fuel cells tend to require much higher operating temperatures (600°C or higher). One of the goals for the DMFC project in our lab was to demonstrate practical power densities using a solid oxide electrolyte such as YSZ or yttrium-doped barium zirconate (BYZ) at temperatures closer to the range of a PEM- type DMFC. Even temperatures as high as 200°C - 250°C could potentially open ceramic electrolyte (SOFC-type) DMFCs for portable commercial applications.

In the following, some further background on DMFCs is discussed, followed by a presentation of a low-temperature YSZ DMFC that uses a PtRu alloy for the anode. The fuel cell architecture was comprised of a dense platinum cathode patterned by nanosphere lithography, a polycrystalline yttria-stabilized zirconia (YSZ) electrolyte, and a sputtered porous Pt0.4Ru0.6 anode. Fuel cell behavior and electrochemical impedance spectra (EIS) were measured. Analyses of the data were performed to extract exchange current densities and activation energies for the fuel cell and component processes. We observed maximum power densities of approximately 2 W/cm2 and 480 W/cm2 at 250°C and 450°C, respectively. These results were consistent with the literature for similar cells and relative performance when using methanol. We speculate that anode reactions involving Ru may contribute significantly to the observed cell impedance. Additionally, we suggest that significant cracking of the methanol may occur at the anode even at these low temperatures.

The experimental details and results discussed in this chapter have been submitted in a manuscript currently under review for publication in the Journal of the Electrochemical Society (44).

4.1 Direct Methanol Fuel Cells

Methanol is a portable, cost-effective and widely available liquid fuel that provides high energy density (18 GJ (LHV)/m3) as well as easy delivery and “instant charging” – for example, swapping out a fuel cartridge or refilling a reservoir, in contrast with charging a battery over several minutes or hours. The latter is particularly important for portable or mobile power applications. Direct methanol fuel cells (DMFC) that utilize an acidic solid polymer exchange membrane electrolyte such as NafionTM have been under development for some time (45; 46), but suffer the persistent and inherent problem of methanol crossover that results in reduced cell voltage, fuel utilization, and conversion efficiency, due to the migration of methanol to the cathode (45; 47; 48). Moreover, DMFCs typically rely on internal reformation of methanol at the anode to produce protons that are transported through the polymer exchange membrane to the cathode where they reduce the oxygen from air to form water. The reforming reaction requires providing to methanol, or carrying along in the DMFC device, at least an equimolar quantity of water. Usually, only dilute methanol concentrations around 1M are employed at the anode compartment to minimize the driving force for methanol crossover to the cathode side. The combined effect of these inherent issues and problems significantly reduces the conversion efficiency to values around 25% at best.

Recent work in our laboratory (49; 50; 51; 52) and others (53; 54; 55) aimed at developing ultra thin film dense proton conducting ceramic electrolyte membranes (56), such as yttria-doped barium zirconate (BYZ), to eliminate issues related to crossover problem by physically blocking methanol. Haile and co-workers followed a similar approach and employed a proton conducting acid solid electrolyte, CsH2PO4, in conjunction with a steam reforming catalyst (alumina-supported Cu-Zn alloy) in a methanol fuel cell, and achieved 226 mW/cm2 at 260°C (57). However, either the use of dense BYZ or acid solid electrolytes still require the availability of stoichiometric quantities of water necessary for the methanol reforming reaction at the anode.

Furthermore, BYZ is susceptible to CO2 poisoning, even under atmospheric conditions for temperatures less than 400°C, via the formation of BaCO3 (56).

Alternatively, the use of an oxide ion conducting, as opposed to a proton conducting, ceramic electrolyte membrane offers the opportunity to achieve direct oxidation of methanol at the anode without the need for reformation. The overall fuel cell operation for proton and oxygen conducting DMFCs is outlined in Figure 4.1. Indeed, methanol fuel operation has been demonstrated previously on SOFCs with SDC or gadolinia doped ceria (CGO) electrolytes and YSZ cermet or lanthanum strontium cobalt ferrate (LSCF) ceramic electrodes (58; 59). Similarly, Jiang and Virkar (60) employed methanol and ethanol fuels in a Ni cermet anode-supported solid oxide fuel cell (SOFC) at 650-800oC with and without steam reformation, and demonstrated 0.6W/cm2 at 650oC and 1.3 W/cm2 at 800oC without reformation. The thickness of the oxide ion conducting yttria stabilized zirconia (YSZ) electrolyte they employed in their SOFC button cells was about 10 m, while the cathode was a mixed conducting oxide composite. To put these power densities in context, we note that polymer electrolyte membrane (PEM) type DMFCs have record power densities on the order of a few hundred mW/cm2 at temperatures between 100°C and 200°C (45; 46).

The benefits of using a YSZ electrolyte include a broad range of stable operating conditions, prevention of fuel crossover, and elimination of the need for reformation of alcohol fuels. On the other hand, ceramic electrolytes suffer from poor ion transport at low temperatures and hence, SOFCs are typically operated at temperatures above 650°C. However, as demonstrated by Jiang and Virkar (60), the ohmic loss from the electrolyte can be compensated for with high surface area electrodes.

4.2 Experimental

The purpose of this study is to demonstrate the feasibility of direct methanol operation of a YSZ-based SOFC at temperatures in the range of 250°C to 450°C with a PtRu anode, which has been shown to be an effective electrocatalyst for methanol oxidation (51; 52; 61). Additionally, we attempt to understand the rate limiting processes through analysis of impedance spectra and comparison to hydrogen benchmark tests on both Pt/YSZ/PtRu and Pt/YSZ/Pt cells previously studied in our laboratory (62; 63).

Polycrystalline 8% YSZ (Ceraflex 8Y) substrates with one side polished and a thickness of 100 m were purchased from Marketech International (Port Townsend, WA, USA). The fuel cell electrodes were deposited from a Pt target (99.99%, Kurt J Lesker) and Ru pieces (99.95%, Kurt J Lesker) by DC sputtering. Methanol, nitrogen, and hydrogen were used as-received.

Figure 4.1 – Schematic diagram of DMFC fuel cell with proton conducting and oxide conducting electrolytes. Proton conducting (a) and oxide conducting (b) electrolytes result in different anode and cathode reactions, with the same net reaction. For the proton conducting DMFC, there is a reformation reaction on the anode side: CH3OH + H2O 3H2 + CO2. In the oxide conducting DMFC, the anode reaction mechanism is more complex and likely depends on electrode material.

Using a Langmuir-Blodgett (LB) technique in connection with nanosphere masking as described by Kim et al. (62), a dense and patterned platinum film was deposited by sputtering onto the polished side of the YSZ wafer. This dense patterned cathode was selected principally for its apparent resistance to thermal degradation and microstructural changes, as observed by Kim (62). Sputtered porous Pt films have a tendency to agglomerate over the course of a few hours at the temperatures used in this study, resulting in significant loss of triple phase boundary (TPB) area. The dense material used in the LB masked cathode does not exhibit this behavior (62), thus allowing for a study of the anode without concern of significant cathode morphology changes affecting the results. The rough sides of the YSZ substrates were then coated with a porous film of PtRu by DC co-sputtering for 150s at 50W under 10 Pa argon. The geometric areas of the LB-masked cathode region and the sputtered anode region were approximately 0.32 cm2 and 1 cm2, respectively. The fuel cell architecture used in this work was previously studied with pure Pt anode by Kim et al under H2 fuel, using the same electrolyte material and thickness as in the present work (62). The conductivity of the Ceraflex 8Y pieces using an LB nanosphere masked cathode and porous sputtered Pt anode was previously characterized by Holme et al. for temperatures below 400°C (63).

Methanol was delivered to the cell by bubbling with nitrogen. A gas mixer in the form of an expansion chamber was placed between the methanol reservoir and fuel cell test stage, with the gas lines heated to 70°C to discourage condensation of methanol. In all cases, the anode fuel mass flow rate was adjusted until no additional current was obtained in the current-voltage (I-V) curve behavior to avoid fuel starvation and severe concentration losses at high current. The cathode was exposed to air. Fuel cell performance was measured over the range 250°C to 450°C.

Electrochemical measurements were recorded using a Gamry Instruments (Warminster, PA, USA) FAS2 potentiostat and E-station. A linear sweep voltammetry method was employed for the I-V measurements, and in all cases, the voltage scan rate was lowered until no further changes in the resulting I-V curve were observed. The

EIS data was measured in potentiostatic mode near open circuit conditions, as well as at lower cell potentials (higher overpotentials), with a signal amplitude of 50mVAC to compromise between noise and averaging effects.

For the analysis presented in this study we do not attempt to determine the active electrode areas and used the cathode geometric area of 0.32cm2 to estimate current and power densities. The cell geometry and architecture used allow us to directly compare the results of H2 experiments to those found with Pt/YSZ/Pt cells by Kim (62) and Holme (63).

4.3 Results and Discussion

4.3.1 Electrode Morphology and Composition

Figure 4.2 shows the scanning electron microscopy images of PtRu electrodes as- deposited and after fuel cell testing. The PtRu anode morphology changes only slightly after several hours of testing with minimal loss of triple phase boundary. Similarly, the LB-masked dense Pt cathode exhibited only minimal morphology change within the test period, consistent with the earlier findings of Kim et al. (62).

Figure 4.2 – SEM image of porous sputtered PtRu anode. Before (left) and after (right) FC operation. Also visible are grain boundaries (- - -) of underlying YSZ (~5um grain size). Only slight morphology change is observed compared to sputtered porous Pt.

Analysis of x-ray diffraction (XRD) data indicated the fcc phase for Pt with the Pt(111) peak slightly shifted, which is consistent with incorporation of Ru into the Pt fcc lattice (52; 61). X-ray photoelectron spectroscopy (XPS) depth profiling of the anode after testing showed the stoichiometry of the cosputtered PtRu was approximately Pt0.4Ru0.6. Phase diagrams of Pt-Ru systems typically show solubility of Ru in fcc Pt up to about 62% (64). It is possible that a small amount of unalloyed Ru is present at the anode, although this was not detected by XRD measurements. Additionally, the XPS measurements showed an O content in the electrode that exceeded the amount expected from observed Zr concentration, as shown in Figure 4.3. As the etching progressed, the Pt and Ru content diminished, and the Zr and O content approached the expected stoichiometry for YSZ. This suggests that some oxidation of the anode may have occurred. Using tabulated thermochemical data (65) for RuO2 and the conditions tested in this study, the cell voltage below which Ru oxidation is likely to occur at the anode ranges from ~550mV at 250°C to ~460mV at 450°C. This supports the possibility of Ru oxidation at the anode by fuel cell operation during polarization tests since the voltage scans typically discharged the cell potential down to 100mV. Measurements of O concentration in the anode of cells operated at low potential for varying time may provide further evidence of this phenomenon, but have not been made.

4.3.2 Fuel Cell Behavior

Sample current-voltage and power density curves are shown in Figure 4.4 and Figure 4.5. The apparent straight-line shape of the I-V curves suggests either significant ohmic limitation or large activation losses. The slope changes when the anode environment is altered, such as by changing between methanol and hydrogen fuel, and the approximate overall slope of the seemingly straight-line IV curves is far greater than the expected ohmic component of fuel cell losses of approximately 4.2 k -cm2 at 250°C, suggesting that electrochemical reaction losses are significant. This is supported by the EIS results, discussed below.

The peak power densities occurred in the range of approximately 0.4-0.5V. The maximum power density data is summarized in Figure 4.6. The maximum power densities observed with methanol at 250°C and 450°C were ~2 W/cm2 and 480 W/cm2, respectively. For comparison, the maximum power densities observed with 2 2 H2 were approximately 10 W/cm and 1 mW/cm at 250°C and 450°C, which is slightly lower than the power density observed for fuel cells with the same architecture but employing a pure Pt anode (62; 66). The difference between our results and the previous studies is most likely due to two factors. First, the SiO2 spheres used for LB nanolithography are larger in the present study, leading to a slightly lower cathode triple phase boundary area (62). Second, although the PtRu anode material is not catalytically less active towards the H2 oxidation reaction (HOR) than a pure Pt anode (67), it is possible that oxidation of Ru at the anode may contribute to the lowering of the fuel cell performance, although it is not clear at present.

Figure 4.3 – XPS depth profiling of PtRu anode after testing. Approximate thickness of PtRu electrode is 100nm. Oxygen content in excess of 2x Zr at% supports the possibility of RuO2 or PtRu-Ox formation.

Figure 4.4 – Current-voltage and power density data at 250°C. Direct methanol (red) and hydrogen data (blue). Voltage sweep rate 10 mV/s. Slope of hydrogen data is ~43.5 k -cm2, much greater than the anticipated ohmic component of 4.2 k -cm2, suggesting activation losses are significant. Sample was first tested using H2 and then CH3OH fuels.

(67), it is possible that oxidation of Ru at the anode may contribute to the lowering of the fuel cell performance, although it is not clear at present.

In general, these power densities have no practical value and indeed are much lower than those published by others (57). However, the peak power density obtained with methanol relative to that measured with hydrogen is consistent with pervious work on direct methanol SOFCs (58; 59). It is important to recognize that the YSZ electrolyte is more stable than BYZ or solid acid electrolytes, both of which require electrolyte humidification. The active area (TPB area) is likely considerably lower for the fuel cells in this study than in studies that utilize highly porous or microstructured electrodes, such as Ni-YSZ cermet anodes. Furthermore, while the electrolyte impedance is not the limiting factor in this study, others in our laboratory have observed a marked increase in the exchange current density of Pt/YSZ/Pt cells when ALD YSZ is used (66).

For comparison, methanol reformation was also tested using a mixture of methanol and water, which resulted in lower fuel cell performance compared to methanol alone. This is to be expected since equimolar water is not required for methanol oxidation when an oxide conducting electrolyte is employed, as oxide ions are delivered to the anode by the electrolyte, as shown in Figure 4.1. Indeed, this is consistent with the findings of previous studies (58; 59), and we have elected to focus our discussion on methanol without water.

Figure 4.5 – Current-voltage and power density data at 350°C. Direct methanol (red) and hydrogen data (blue). Voltage sweep rate 10 mV/s. Sample was first tested using H2 and then CH3OH fuels. Same sample as in Figure 4.4.

Figure 4.6 – Peak power density (mW/cm2) plotted on a log scale. Direct methanol results shown in red, with hydrogen data in blue for reference. Data averaged across all viable samples tested.

4.3.3 Methanol Cracking

It is worth mentioning the possibility of methanol cracking at the temperatures tested. Using simple estimates from tabulated thermochemical data (65), one can show that the reaction

CH 3OH CO 2H2 4.1 becomes energetically favorable at approximately 137°C, which is much lower than the temperature range examined in this study.

Thermodynamics does not provide conclusive evidence that methanol cracking occurs, and it is unclear whether the overall methanol oxidation reaction involves the oxidation of adsorbed methanol, or includes methanol cracking products of Equation 4.1. However, it is possible that cracking of methanol may occur to a certain extent at the temperatures in this study, thereby effectively providing a mixture of H2, CO, and

CH3OH to the anode surface. Previous studies also suggest that methanol cracking may play a significant role at the anode (58; 59; 60).

At high temperatures, we might expect more methanol cracking, and therefore less difference between CH3OH and H2 operation of the fuel cell, which is consistent with the converging impedances and cell performance at high temperatures. However, at this time, further exploration of the mechanisms of YSZ-supported PtRu catalysis of the hydrogen oxidation reaction (HOR) and methanol oxidation reaction (MOR) is needed to fully explain the EIS results, and gas analysis of the anode inlet and outlet streams may help elucidate the extent to which methanol decomposes before being oxidized at the anode.

4.3.4 Fuel Cell Exchange Current Density

The activation overvoltage was determined by

act total ohm (Vmeas,OC Vmeas ) iRion 4.2 where Vmeas,OC is the measured voltage at open circuit conditions, and iRion is the ohmic overpotential due to the ionic resistance of the electrolyte. Exchange current

65 density, j0, and transfer coefficient, , were estimated through fitting of the ohmically- corrected IV data using the high overpotential Tafel approximation of the Butler- Volmer equation:

F ln j ln j 0 RT act 4.3 where j is the measured current density. The slope of this fit can thus be related to the transfer coefficient by = (slope) x RT/F, where R is the ideal gas constant, and F is

Faraday‟s constant. The intercept of the linear fit is ln j0. The subscript on overpotential, act, is used to indicate that we are only fitting the overpotential to kinetic considerations; we assume no mass transport losses, and have corrected for charge transport in Equation 4.2. The Butler-Volmer equation (and thus the Tafel approximation) assumes a single step one-electron process (5). Therefore, the analysis presented here results in an “overall” exchange current density and transfer coefficient for the fuel cell. This approximation is often sufficient for H2/air fuel cells, which are typically limited by the oxygen reduction reaction (ORR) (4).

It should be noted that there is some discrepancy in the literature as to whether or not a factor n should be included in the exponent of the Butler-Volmer equation and therefore in the slope of the Tafel approximation. The factor n represents the number of electrons transferred in the rate limiting step. We follow the assumption that the rate limiting step is most likely a single electron transfer and that the factor n is therefore unity and may be dropped from the expression (5). Therefore, we have chosen to report the transfer coefficient for n = 1 and leave any investigation of the reaction mechanisms to future studies. This may produce some error in the reported value of the transfer coefficient if there are multiple reactions that should be considered. As such the values of reported here should be considered fitting parameters rather than the true transfer coefficient for the rate-limiting reaction.

Figure 4.7 shows the kinetic parameters as determined by the above method. The activation energy was determined from the slope of a linear fit of the natural logarithm of exchange current density plotted against reciprocal temperature. The overall

66 activation energy for the fuel cells under direct methanol was found to be 0.90 eV. We found the activation energy for hydrogen operation to be 0.81 eV.

Figure 4.7 – Kinetic parameters determined by Tafel fitting. (a) Natural log of exchange current density plotted against reciprocal temperature. (b)Transfer coefficient. Methanol results shown in red, with hydrogen results shown in blue for reference.

If the oxygen reduction reaction (ORR) were the sole limiting reaction, as is often assumed for H2/air fuel cells (4), we would expect to observe no change in the exchange current density activation energy between methanol and hydrogen anode environments. Therefore, it is likely that some anode process or processes, in addition to the cathode ORR, act to limit the fuel cell performance. This is supported by, and further discussed with, the EIS measurements.

The transfer coefficient remained close to 0.10 for methanol and hydrogen fuel across the temperature range investigated, indicative of a high degree of asymmetry in the free energy space of the rate limiting step.

4.3.5 Electrochemical Impedance

Figure 4.8 shows typical EIS data and a sample fit. In the discussion that follows, the three easily identifiable impedance loops will be referred to as loops “1”, “2”, and “3”, in order of decreasing frequency, as indicated in Figure 4.9. To gain better insight about the processes they represent, we compare the H2 data to that found by Holme (63) for the Pt/YSZ/Pt cell with LB-masked cathode (“NSL cathode” sample in (63)) and discuss the impedances measured using methanol in the context of that comparison. Holme found three distinct impedance loops in the Nyquist plot and, in conjunction with Monte Carlo simulations to estimate reaction rate for several possible reactions, suggested that the highest frequency loop was due to vacancy diffusion in the electrolyte, the intermediate frequency loop was due principally to chemical reactions at the anode, and the lowest frequency loop was related to cathode chemical and electrochemical reactions as well as water formation on the anode. We speculate that the impedance loops described in (63) correspond to the loops observed with H2 fuel in the present study as follows:

Loop “E” – “HF” in Holme; electrolyte impedance, calculated with R fixed during fitting.

Loop “1” – no equivalent in Holme; likely related to anode processes involving Ru

Loop “2” – “IF” in Holme; suggested to be H2(g) → 2Hads or Oads + Hads →

OHads.

Loop “3” – “LF” in Holme; suggested reactions include O2(g) → 2Oads and anode water formation.

Figure 4.8 – Sample EIS data taken at 300°C. a) H2 fuel (x), and b) methanol fuel (○). H2 data shown again in methanol plot as for reference. Corresponding 4 R||C equivalent circuit fit (solid line) for methanol near open circuit conditions using electrolyte conductivity data from Holme (63). Cathode geometric area 0.32cm2, anode geometric area 1cm2. Bias conditions near open circuit (black) and at +0.4V vs anode (red). Loops marked as “1”, “2”, and “3” for discussion purposes. Expected electrolyte impedance ~2.3 k (Loop “E”, not indicated in this figure).

Figure 4.9 – Sample EIS data taken at 250°C. H2 fuel at the anode near OCV conditions. Solid markers are for saturated fuel flow rate (40 sccm), hollow markers are for low fuel flow rate (5 sccm). The highest-frequency loop appears to obscure a smaller loop near the high-frequency intercept that does not vary in size when the fuel flow rate is decreased. The expected electrolyte impedance is roughly 13 k at this temperature. Solid line loops sketched in the inlay illustrate the possibility of two processes merging in the high frequency range – electrolyte loop “E”, and highest frequency reaction loop “1”.

It is expected that the geometric capacitance of the electrolyte is much smaller than the double layer capacitances at the electrodes. Therefore, the highest frequency loop is typically attributed to the electrolyte ohmic resistance coupled with electrolyte geometric capacitance. However, loops 1, 2, and 3 in the present study all increase in magnitude when fuel flow rate is reduced, as in Figure 4.8, or the fuel is changed from

H2 to methanol, as in Figure 4.9. This observation suggests loops 1, 2, and 3 are all related to electrode processes, since the electrolyte impedance should be constant with all variables except temperature. The electrolyte impedance we expect to observe, based on the results for the LB-masked Pt/YSZ/Pt cells presented by Holme (63), is smaller than the impedance of loop 1 and larger than the high-frequency intercept with

70 the real axis of impedance, except when the electrolyte characteristic frequency surpassed 300 kHz, which is the measurement upper limit of our equipment. Furthermore, loop 1 exhibits a slight shoulder near the high frequency intercept, as seen in Figure 4.8 and Figure 4.9. We observed this feature in both methanol and hydrogen at 250°C and 300°C, and therefore suggest that the electrolyte impedance (loop E) is obscured by the larger loop. The general shape of the Nyquist plot is similar between the fuels used, suggesting similar processes, and examination of impedance loop activation energies will help characterize differences.

To fit the EIS data we utilized an equivalent circuit model comprised of a total of four R||C loops in series, with the electrolyte impedance calculated from the cell geometry and known conductivity data for the YSZ substrates (63). For the measurements at 350°C and 450°C, the characteristic frequency of the electrolyte impedance loop (loop E) was most likely higher than the frequency upper limit of the instrument. In these cases, the equivalent circuit was a resistor, representing the electrolyte, in series with three R||C loops for the electrode processes. The capacitive elements used were constant phase elements (CPEs) rather than ideal capacitors to improve the quality of the fit. It has been suggested that CPE behavior may be due to surface roughness (68) or local variations in the electrode material that may result in a distribution of activation energies for a reaction (69).

Aside from the presence of the additional large high frequency impedance (loop 1), the Pt/YSZ/PtRu cells behave similarly to the Pt/YSZ/Pt cells from prior studies when operated under H2 fuel at the anode. This suggests there may be additional processes due to the change in anode material from Pt to PtRu. The resistive element fit values for loops 1, 2, and 3 are plotted in Figure 4.10. The activation energies were calculated for each bias when the resistances were plotted against reciprocal temperature, and are catalogued in Table 4.1. The capacitance associated with each loop was calculated using the peak frequency measured for each loop (i.e., where -ZIm is at a maximum) and the resistance from the fit using the relationship 2 f = 1/RC.

Figure 4.10 – Typical results from fitting the EIS data to four R||C loops. The resistor and characteristic frequency values are shown for the three non-electrolyte loops. Direct methanol fit results are shown in red, hydrogen in blue. When bias dependence was observed, OCV ( ), 0.7V v anode ( ), and 0.4V v anode (○). Linear fits of resistance are shown as solid lines, with activation energies listed in eV and also summarized in Table 4.1 (inset).

Table 4.1 – EIS loop activation energies (eV).

Loop CH3OH H2 Bias 1 1.11 0.97 (indep)

2 1.23 0.985 (indep)

0.58 0.36 OCV 3 0.78 - 0.7V vs. anode 0.82 0.62 0.4V vs. anode

Figure 4.11 – Area specific capacitances of four R||C loop model. Calculated from resistances determined by EIS fitting and frequencies measured from raw data. Area 0.32 cm2. Methanol (top) and hydrogen fuel (bottom). Loop 1 (black), loop 2 (blue), loop 3 (red). When bias dependence was observed, OCV ( ), 0.7V v anode ( ), and 0.4V v anode (○).

Loop 1 represents a large impedance with high characteristic frequency that is weakly, if at all, dependent on cell potential, but exhibits significant dependence on the fuel delivered to the cell, distinguishing it from electrolyte impedance. We suspect, based on the absence of this loop for Pt/YSZ/Pt cells, that the Ru plays a significant role in this impedance. The difference in activation energy for loop 1 between methanol (1.11 eV) and H2 (0.97 eV) reinforces the notion that loop 1 is related to anode processes. At present the specific process is unclear, but possible reactions are Ru + O → RuOx , Ru-

Hads + O → Ru-OHads and Ru-OHads + O → Ru-H2Oads. We speculate that it is unlikely that loop 1 is directly connected to carbon processes since it is also present when the fuel stream contains only pure hydrogen. As seen in Figure 4.11, the specific capacitance for loop 1 ( < 1 nF/cm2) is significantly lower than the usual double-layer regime of 10 F/cm2. The physical origin of this capacitance and its associated loop is not understood at this time.

For loop 2, we observed activation energy of 0.98 eV under H2. Holme observed approximately 1.01 eV for the LB masked cathode-side loop (i.e., “IF” loop) (63). We speculate that the similarity in activation energies for hydrogen suggests that the IF loop in Holme represents the same anode process(es) as loop 2 in the present study. We found significantly larger activation energy, 1.23 eV, for loop 2 when methanol fuel was employed. Additionally, we observed almost no bias dependence in the impedance of loop 2, consistent with the previous study on Pt/YSZ/Pt cells. These observations further support the hypothesis that loop 2 in the present study and the IF loop in Holme represent the same anode processes under H2 fueling. Possible reactions include H2 adsorption and oxidation of adsorbed H. Kinetic monte carlo simulations have shown these reactions to be faster than ORR processes and other HOR steps, but slower than YSZ oxygen vacancy processes (63).

The bias independence of loops 1 and 2 suggest that the processes are fully activated at low overpotentials, and a large fraction of overpotential is discharged through other reactions, namely, those exhibiting strong bias dependence as in loop 3.

For loop 3, we found the activation energy to increase significantly with decreasing cell voltage. This is consistent with the findings for the “LF” loop in Holme (63), suggesting some similarity in processes. However, the activation energies are significantly larger in this study for both H2 and methanol fuel. We speculate that loop 3 may include anode processes more severely affected by the use of PtRu instead of Pt than the processes in loop 2. The significant bias dependence suggests a sluggish process, such as ORR at the cathode. However, the increase in activation energy when methanol fuel is used suggests anode processes may be merged with the loop as well. The anode processes may be reaction between adsorbed H and OH, water evolution, or a slow step in the transition of oxide ions from YSZ to the anode.

The higher activation energies (Table 4.1) for methanol are neither surprising nor unexpected, given that the anode reaction pathway is likely more complicated than for the case of hydrogen fuel. This is consistent with the findings of Liu et al. (59), although the values are somewhat different due to the differences between fuel cell architectures used in the two studies. We speculate that each impedance loop observed with methanol fuel represent processes on the same electrode as in the H2 case. That is, under methanol, loop 2 represents anode processes, and loop 3 mainly represents cathode losses and possibly anode water formation.

It is not clear from the measurements made in this study which impedance loop, if any, represents CO oxidation processes or includes steps of the “bifunctional mechanism” of Pt-COads oxidation. It is also not evident whether the bifunctional mechanism occurs as described in the literature, in which methanol is delivered as aqueous solution (70; 71).

In general, the literature does not contain much information regarding the HOR or vapor-phase MOR on PtRu alloys, particularly for oxide conducting electrolytes. It is established, however, that the MOR mechanism is complicated and there are several possible pathways that may take place. Some of the possible reaction steps are outlined in previous studies (58; 59; 60; 72). Additionally, a combination of PtRu catalyst, oxide conducting electrolyte, and vapor phase methanol is uncommon. This

75 makes speculation of specific electrode processes represented by the impedance loops difficult. Quantum chemical calculations and anode product gas analysis may help elucidate the MOR mechanisms in future studies.

4.4 Conclusions

We have demonstrated direct methanol operation of a Pt/YSZ/PtRu solid oxide fuel cell at temperatures between 250°C and 450°C. It is at present unclear as to whether this is truly a direct methanol oxidation or if methanol cracking plays a significant role in cell performance. However, based on the similarities between impedance curves of

H2 at different flow rates and methanol fueling, we speculate that significant cracking may occur, resulting in a mixture of CH3OH, CO, and H2 at the anode surface, with the additional reactions of CO oxidation and uncracked CH3OH oxidation contributing to the impedances and higher activation energies.

Electrochemical impedance spectra show a large high-frequency Nyquist loop that is not associated with the electrolyte impedance, and may include oxidation of the PtRu anode and chemical reactions at the anode. The lower frequency impedance loops are attributed to anode and cathode electrochemical reactions and are consistent with previous findings for H2 fuel using a pure Pt anode. We speculate that anode processes contribute significantly to all of the observed electrochemical (i.e., non-electrolyte) impedance loops, as evidenced by the increase in activation energy when methanol is used.

There are still unanswered questions regarding the anodic behavior of the Pt/YSZ/PtRu fuel cell, and future studies would benefit from examination of the composition of anode inlet and outlet streams.

As expected, the YSZ-based DMFC produced lower power densities than the same fuel cell supplied with a stream of pure H2 at the anode. Higher power densities may be obtainable using pure methanol vapor (rather than methanol in N2) and ALD thin film YSZ electrolytes with high-surface area electrodes.

5 High surface area electrodes

This chapter explores the fabrication work and experimental results for SOFCs with high surface area anodes. A brief review of the motivation behind using a high surface area electrode is followed by a discussion of mixed conducting materials. The bulk of the chapter will focus on efforts towards the fabrication of electrodeposited nanowire fuel cell electrodes. Lastly, some alternate strategies to nanowire electrode preparation are presented, including the use of a liquid-crystal phase template.

5.1 Increasing j0 in fuel cells

The reaction kinetics and activation losses for fuel cells were discussed in §2.3 and §0, respectively. From the Butler-Volmer equation and the Tafel approximation

(Equations 2.12 and 2.17), it is clear that a higher exchange current density, j0, yields a higher net current density for a given overpotential (or lower overpotential for a given net current density), resulting in better fuel cell performance and higher maximum power density. The exchange current density can be expressed as

Ea j0 nFcf exp 5.1 kBT where c is reactant concentration, f is a temperature-dependent rate constant, kB is the

Boltzmann constant, and Ea is the activation energy barrier to reaction. From this expression, three means of increasing the exchange current density are obvious: increase reactant concentration (or gas pressure), increase the temperature at which the reaction occurs, or lower the activation barrier through catalysis of the reaction (such as with the surface treatments in Chapter 3). Additionally, and not as immediately apparent, j0 can be increased by increasing the number of reaction sites available. Current density in a fuel cell is determined from the geometric (i.e., planar) area rather than the real reaction surface area. Therefore, increasing the surface area available for the reaction results in a greater current per unit geometric area:

j0' j0 Areal Ageom 5.2

However, for an increase in exchange current density to have a significant impact on the fuel cell performance, an orders-of-magnitude change is needed due to the logarithmic relationship between activation overpotential and current density, as seen in the Tafel approximation. Figure 5.1 illustrates this point for a hypothetical fuel cell reaction with three different exchange current densities and all other parameters the same. In the example in Figure 5.1, the power density in the high current density regime is only about 6% greater when the exchange current density is doubled, but roughly 20% greater when the exchange current density is increased by a factor of ten.

Figure 5.1 – Example activation losses for different values of j0. 2 2 Calculated from the Butler-Volmer equation for three values of j0 : 1 A/cm (blue), 2 A/cm (red), 10 A/cm2 (green). Other parameters used are OCV = 1.2 V, = 0.5, T = 300 K. Power density at j = 10 mA/cm2: 6.05 mW/cm2 (blue), 6.41 mW/cm2 (red), and 7.24 mW/cm2 (green). Note that this model ignores ohmic and concentration losses.

One means of achieving a high reaction area is to structure the reaction surface in three dimensions. Nature accomplishes this through folding of membranes as seen in mitochondria and chloroplasts (3). A previous study from our lab was able to demonstrate improved fuel cell performance using a “corrugated” fuel cell structure in which the electrode/electrolyte membrane was fabricated as an array of cup-like structures. The total fuel cell membrane area was increased by a factor of

78 approximately four relative to a planar thin film fuel cell fabrication scheme, and the maximum power density achieved increased by a factor of two with all test conditions equal (73).

Figure 5.2 shows a SEM image of the corrugated thin-film fuel cell structure and the resulting I-V and power density behavior. These results confirm the notion that by structuring the fuel cell to mimic the folds found in naturally occurring reaction surfaces, a greater reaction rate per device geometric area can be attained.

The goal of the work presented in this chapter is to explore the effects on fuel cell performance of structuring only the electrodes. That is, we are studying the effects of selectively increasing only one fuel cell interface active surface area instead of increasing all of them (cathode/oxidant, cathode/electrolyte, electrolyte/anode, anode/fuel). Furthermore, it may be possible using certain techniques to increase the reaction surface by multiple orders of magnitude. Depending on fabrication processes, it may also be possible to combine the techniques discussed in this chapter with the corrugated structure studied previously by Su (73).

Figure 5.2 – Fuel cell behavior for corrugated thin film structure. I-V and power density behavior (left) for corrugated and flat thin film YSZ fuel cells (SEM, right). Cup diameter and depth 15 m and 20 m, respectively (73).

5.2 Mixed electronic ionic conducting (MEIC) electrodes.

As discussed in Chapter 2, fuel cell reactions require a triple phase boundary (TPB) to occur. Therefore it is important to maintain TPB density on the electrode surface during structuring, limiting the materials options for structuring by requiring that as much of the structured electrode is accessible to ions, electrons, and reactants. For example, it would not be appropriate to use dense platinum for the electrode shown in Figure 5.3. In addition to the material cost, such a device would likely not have sufficient TPB interface for improved operation. However, if the material used for the electrode in Figure 5.3 were able to conduct both electrons and the ion species conducted by the electrolyte, the fuel cell would be viable, particularly so if the surface of the electrode were sufficiently catalytically active towards the reactions at the interface with fuel or oxidant.

Figure 5.3 – Schematic diagram of SOFC with one structured electrode. The bottom electrode is porous to allow the gas phase to reach the electrode/electrolyte interface. In order for the structuring of the top electrode to result in a functioning fuel cell, the material must either support both electrons and ions, or be porous like the bottom electrode. The former condition results in the entirety of the surface functioning as triple phase boundary.

In fact, such materials do exist, and are referred to as mixed electronic ionic conducting (MEIC) materials. The advantage to MEIC electrodes is that the entire surface of the material is potential TPB area provided catalyst is uniformly deposited, as is achievable with atomic layer deposition. By nanostructuring MEIC electrodes, very high TPB areas are achievable. Examples of MEIC materials are lanthanum strontium cobalt ferrite (LSCF), an oxide conductor used in some solid oxide fuel cell

(SOFC) cathodes, and Pd (74) and RuO2 (75), both of which are compatible with proton conducting electrolytes, such as yttrium-doped barium zirconate (BYZ), a principal material of interest in the DMFC studies in our lab. In the case of a catalytic material capable of conducting electrons and ions, such as Pd, the two phases are collapsed into one mixed conducting material, but we still refer to TPB, even though the TPB consists of a single interface (electrode/fuel).

In the case of Pd, the “proton conduction” is more accurately described as diffusion of absorbed atomic hydrogen. Hydrogen atoms occupy the octahedral interstices of face– centered cubic (fcc) Pd. Thus the maximum theoretical stoichiometry of Pd hydride is near Pd1H1, as there are four octahedral interstitial sites (body center and edge centers) within the fcc unit cell (containing four atoms). The ordered-phase limit of H absorption by Pd is closer to PdH0.6 at room temperature (74). Absorption of H by Pd results in some swelling of the material, as well as a conductivity change. Therefore, the Pd-H system is of interest for H2 sensing applications. (76; 77)

Because of the nature of H absorption and transport in Pd, it is more appropriate to consider the diffusivity of H in Pd than the conductivity of H+. The literature contains several reports regarding the diffusion characteristics of H in Pd. A survey of several studies (78) yields a temperature-dependent diffusivity estimate of

2 3 cm 0.24eV DH in Pd 4.8 10 exp 5.3 s kBT At 350°C, and a diffusion length equal to the entire thickness of a 50 m Pd foil, the limiting current density (Equation 2.22, for PdH0.5 at surface and H depletion at electrolyte interface) is approximately 50 A/cm2. This is far beyond any current densities observed on the fuel cells studied in our lab. Indeed, characterization of fuel cells using Pd foil substrate/anode, BYZ electrolyte prepared by pulsed laser deposition (PLD), and sputtered porous Pt cathode and anode catalyst showed no significant difference in fuel cell behavior between 25 m and 50 m Pd foils. The 25 m Pd foils were more difficult to work with as they were easily bent and torn during routine processing and testing. These tests demonstrated the feasibility of Pd-foil

81 based BYZ SOFCs, and confirm that the Pd thickness is not the limiting factor in performance.

This work studies Pd for electrode nanostructuring on BYZ electrolytes. While RuO2 would also have been appropriate, Pd was selected for the initial studies presented here due to the simpler deposition methods for Pd (including DC sputtering) and availability of Pd foils. As a consequence of this selection, the experiments are limited to fuel cells with structured anodes out of concern that cathode-side Pd may suffer damage from oxidation at elevated temperatures. RuO2 is stable in air at the temperatures studied (65) and may be suitable for cathode structuring. In the following sections, a few different Pd structures are studied, beginning with nanowires.

5.3 Pd Nanowires

The first high surface area structure investigated was nanowires. Nanowires are particularly attractive due to the potential for a high aspect ratio and thus large surface area per amount of material used. To illustrate this point, consider the geometry shown in Figure 5.4. The regular hexagonal array of cylindrical nanowires has a surface roughness (real area normalized by geometric area) given by

A 2 d h real 1 2 5.4 Ageom 3 D where d is the wire diameter, h the height, and D the center-to-center spacing between wires. For wires grown in our lab, a roughness of 5 to 10 per micron of wire height is typical. It is not clear at this time how much of the nanowire length contributes to increased current density. That is, we expect diminishing returns with added nanowire length beyond a certain point due to hydrogen mobility in the wires and gas diffusion. It is possible that surface diffusion of H on Pd will be faster and the hydrogen diffusion along the wires will not limit the device performance. This section will discuss the issues encountered in nanostructuring high surface area Pd anodes for BYZ-based fuel cells.

Figure 5.4 – Schematic of close-packed nanowire geometry. Each wire occupies a hexagonal area of the substrate surface. The total surface area is equal to the planar area plus the area of the nanowire walls, as expressed in Equation 5.4.

5.3.1 Pd Nanowire Growth and Characterization

Palladium nanowire growth was carried out electrochemically. A porous membrane template was coated on one side with approximately 100 nm of Pd with a thin Ti adhesion layer by DC sputtering. Constant potential or constant current conditions were imposed between the Pd/template substrate and a Pt or Pd counter electrode. The electrolyte used was initially 5 wt% PdCl2 in 10 wt% HCl. We found that constant potential of –0.3V v AgCl/Ag (3M KCl) resulted in a nanowire growth rate of roughly 200nm/s for a template with 80nm diameter pores and roughly 2x109 pores/cm2. By diluting the concentration to 0.5 wt% PdCl2, we were able to slow growth by roughly an order of magnitude under the same conditions. Constant current deposition was carried out with typical current densities of less than 10 mA/cm2. Wires deposited by constant current have not been fully characterized in our lab, but have not shown significant differences in behavior. Unless otherwise indicated, the analyses presented are for constant potential depositions.

Two types of commercially available template membranes were used. The first was a thin sheet of track-etched polycarbonate (TEPC). The second was anodized aluminum oxide (AAO). TEPC templates with pore diameters on the order of 100 to 200nm were purchased from various sources, and under certain conditions, growth resulted in apparently hollow nanotube-like depositions, shown in Figure 5.5. We did not

83 determine the cause of the hollow deposition nor fully characterize the parameters for growing wires versus tubes, but suspect that certain current densities (i.e., deposition rates) favor growth along the walls of the TEPC pores. This behavior was also observed with RuCl3 deposition solutions before Pd was selected as the sole focus of the nanowire efforts. Fukunaka et al. (79) report conditions for Ni, Co, and Fe nanotubes using TEPC templates.

The AAO templates were purchased from Synkera, Inc, and had pore diameters ranging from 35nm to 100nm. The pores in AAO are parallel and evenly spaced as a result of the membrane preparation process, whereas the TEPC pores show some randomness in spacing and are not completely parallel. Formation of tube-like wires was not observed with AAO templates.

After wire growth, TEPC templates were removed by a wash of dichloromethane

(CH2Cl2), and AAO templates were removed by soaking in 1 M phosphoric acid at 60°C until the alumina was completely dissolved, typically 2-3 hours. Figure 5.6 shows Pd wires ~20 m long in the 50 m AAO template and the resulting structure after removal of the AAO template.

Figure 5.5 – TEPC-templated nanotube-like structures. Pd (left) and RuOx (right). The sheet of material in the upper left of the Pd image is likely a piece of the Pd film sputtered on the back of the TEPC, and shows the distribution of pores in the TEPC, supporting the idea that growth was favored along the walls of the template rather than in the center.

Figure 5.6 – AAO template Pd nanowires. Cross-section (left) with embedded Pd wires (strip of brighter contrast, circled); bright stripes of contrast above wires are pore edges. Not visible is a sputtered thin film of Pd used as the cathode for electrodeposition. Thin-film supported Pd wires after removal of the template (right).

Figure 5.7 – Comparison of AAO-templated Pd wire rinsing methods. a) Pd nanowires grown in an AAO template after H3PO4 etching with only water rinse, left, and (b) shorter wires with stepped ethanol rinses, right. The rough texture of the thin film Pd electroplating cathode is visible in both images above, a result of sputtering onto AAO.

In order to better study the effects of nanowire length on fuel cell performance, we opted to reduce the deposition time (and therefore length of wires), thus avoiding the “tangled noodle” morphology seen in Figure 5.6. Despite this measure, wires had a tendency to agglomerate. We found that by implementing a short growth period and a stepped ethanol rinse after removal of the AAO, wire agglomeration was practically eliminated, as shown in Figure 5.7.

In order to determine the real roughness factor due to the nanowires in Figure 5.7a, we measured Pd characteristic cyclic voltammograms (CVs) in 0.5 M H2SO4 (aq). A clean Pd foil was used as the substrate for the thin-film supported nanowires. The Pd film-supported wires were attached to the foil by taping the AAO template to the Pd foil. Prior to the CV tests, both the nanowire sample and a clean Pd foil piece were coated with roughly 4 nm of Pt by atomic layer deposition (ALD). This ALD Pt layer was used because the end SOFC device design included a uniform coating of Pt as catalyst on the structured electrode. Figure 5.8 shows the CV results. The roughness was determined by the oxide stripping peak height, which scales linearly with the real surface area in linear sweep and cyclic voltammetry.

Figure 5.8 – Characteristic CVs of Pd film and Pd film with wires. Cyclic voltammograms measured in 0.5 M H2SO4 for sputtered Pd film (left) and sputtered Pd film with wires (right). Note difference in scale on vertical axes. Samples were coated with ~4nm Pt by ALD, as would be done for a sample used in fuel cell testing. Based on length and spacing of wires, roughness of ~28 is expected. Observed value from CV is ~14. SEM of sample is that shown in Figure 5.7a.

The measured roughness was approximately half of the estimate from Equation 5.4 for the diameter and density of nanowires, suggesting that the degree of agglomeration may be an important factor. The wires in Figure 5.7a appear to have a lower yield (number of pores filled during plating) than those in Figure 5.7b, another possible source of the discrepancy between measured and predicted surface area, likely due to surface contamination prior to wire growth.

Figure 5.9 – TEM analysis of Pd nanowire. Micrograph and observed (111) lattice spacing and diffraction pattern. TEM images by Munekazu Motoyama.

Figure 5.9 shows the results of TEM measurements made to determine the phase and crystallinity of the wires. The diffraction pattern is consistent with polycrystalline fcc Pd, and the lattice spacing is consistent with the (111) spacing in Pd; we estimate 2.36 Å from the image, which is within 5% of 2.25 Å calculated from the lattice parameter for fcc Pd. The micrograph in Figure 5.9 shows a polycrystalline wire with a diameter of approximately 80 nm.

The above section demonstrates that we were able to achieve growth of Pd nanowires, controlling the diameter by selecting appropriate templates, and control length by limiting deposition time and electrolyte concentration. In the next section, we present a discussion of the issues of joining the wires with a viable substrate for fuel cell fabrication and various strategies for managing these issues.

5.3.2 Pd Nanowire Fuel Cell Fabrication Processes

The most critical issue encountered during our nanowire studies was that of incorporating the nanowire growth with a viable fuel cell substrate. In this case, “viable” simply means a substrate robust enough to be handled during testing and transport around the laboratory, as well as chemically compatible with any further processing required for the nanowire electrode. Several attempts were made, most of which did not result in testable fuel cells. The methods used are presented below, followed by a brief discussion of potential solutions to the problem.

5.3.2.1 Dissolution of template on substrate

The first method attempted to attach the thin-film supported wires to a silicon or Pd foil substrate by dissolution of the nanowire templates while the sputtered film side was in contact with the substrate. The results of this method are highly variable, and seem largely dependent on the direction of fluid motion of the solution used for template removal. In the case of TEPC, template dissolution is very rapid and the thin film supported nanowire structure is prone to tearing and folding, resulting in significant loss of nanowire area. While the AAO is more slowly etched, removal of samples from etchant tended to result in the same sort of tearing and folding seen with

TEPC. Relying on this sort of lamination process to join the thin film supported wires and underlying substrate material may increase the contact resistance between the materials, potentially negatively impacting fuel cell performance by increasing ohm. The samples in Figure 5.5, Figure 5.6, and Figure 5.7 were prepared by this method. Numerous tears made it extremely difficult to prevent the thin film Pd from folding over itself or crumpling, resulting in reduced coverage of the substrate. This problem may be limited by tightly controlling the fluid motion of the chemicals used to remove the templates, however, we were not able to demonstrate consistent transfer of the thin film supported wires to any substrate by attaching the template to a substrate and removing the template.

5.3.2.2 Local template etching

By etching only a small region of the rigid AAO template, a region of Pd could be exposed, making a functional fuel cell, as shown in Figure 5.10 below, where the AAO not etched acts as a support structure for the exposed thin film fuel cell.

Figure 5.10 – Schematic diagram of a BYZ-based fuel cell built on locally etched AAO membrane.

The primary problem with this method is that liquid etchant must be used, as Pd is more susceptible to dry etching than Al2O3. It is difficult to etch only a small area using isotropic wet etchants, resulting in a relatively large area of free standing thin film Pd/BYZ. The devices fabricated in our lab (66; 50) limit the free-standing membrane area to a few 100 m on a side. Using this method we were unable to demonstrate through-etching of the AAO without destroying the thin film Pd membrane as well.

5.3.2.3 Joining to substrate by electrodeposition

Another route explored for joining nanowires to a viable substrate was to “weld” the Pd thin film-coated template and a substrate by electrodeposition of Pd or other metals. A flat, clean substrate, either metal foil or a silicon wafer piece, was prepared with a through-hole approximately 2 mm in diameter. In the case of a silicon wafer, the back (unpolished) side of the wafer piece was coated with a thin film of gold by DC sputtering. The metal foils used were most commonly Ni, Cu, or Pd. An AAO template with Pd wires plated into the pore channels was taped to the front side of the substrate, covering the through-hole. A film of Pd was then electroplated on the backside with a thickness of about one micron. After electroplating, the samples were rinsed in water and the tape was carefully removed with tweezers and soaking in ethanol. The samples were then immersed in phosphoric acid to remove the AAO template. Figure 5.11 shows an SEM image of the backside of a sample after this process and a schematic diagram of a fuel cell that could be fabricated using this strategy. As with many of the other methods discussed, the outcome is extremely sensitive to fluid motion during processing resulting in tears in the Pd film.

Figure 5.11 – Thin-film supported nanowires joined to substrate by electroplating. The SEM image shows the poor quality of the electroplated Pd “weld”. The schematic shows thin-film electrolyte and cathode deposited on the back side of the substrate, visible in the SEM.

5.3.2.4 Thick electroplated Pd substrate

By electroplating a thick layer of Pd on commercial AAO, a suitable substrate may be formed. We have deposited thick Pd from PdCl2 in HCl, as shown in Figure 5.12. It is known that PdCl2 is a suitable chemistry for electroplating thick deposits of dense Pd (80). However, significant cracking was observed in our samples. Therefore, this method may be acceptable if the Pd growth conditions are adjusted to reduce the internal stress of the deposit. Once wires are plated in the AAO and the thick Pd is formed, the AAO may be removed and BYZ deposited on the flat side, followed by catalyst deposition.

Figure 5.12 – SEM image of commercial AAO with 6.5 m of Pd deposited by electroplating.

5.3.2.5 Drop casting and annealing

By growing wires in a commercial template (either TEPC or AAO), then removing the template and depositing the pre-formed wires from solution onto a Pd fuel cell anode, a high surface area MEIC anode may be formed. It is possible that a brief electrodeposition of Pd or thermal annealing may be necessary to ensure strong adhesion between the wires and the substrate. We have also considered that the incorporation of a fracation of nanowires made of magnetic material, such as Ni, may

91 help orient the wires in a parallel array if deposited onto the surface under a magnetic field. At present, this is speculation and we have not yet tested this strategy.

5.4 Growth of Template on Substrate

In an effort to work around the problems of attaching nanowires to a substrate after growth discussed above, we attempted to form AAO templates directly on silicon wafer and Pd foil substrates. We were also able to fabricate a free-standing membrane of AAO/Pd/BYZ on a silicon wafer, but were unsuccessful in uniformly plating wires into our AAO membranes. It is not practical to prepare a TEPC template on a silicon wafer due to the use of heavy ion bombardment in track formation prior to etching.

5.4.1 Aluminum anodization

Anodization of Al to Al2O3 is well understood and there are several anodization baths used in fabrication (81). In the anodization process, a high voltage is applied between an aluminum work piece and a counter electrode, with a dilute acidic electrolyte. Sulfuric, oxalic, and chromic acid are the most commonly used electrolytes, with temperature and voltage conditions varying with solution. The aluminum surface (anode) is oxidized by the water in the electrolyte solution, and the aluminum oxide is simultaneously dissolved by the protons in the electrolyte:

+ - 2Al + 3H 2O Al 2O3 + 6H + 6e 5.5

+ 3+ Al 2O3 6H 2Al + 2H 2O 5.6 The cathode reaction is simply the formation of molecular hydrogen:

- 6H 6e 3H2 5.7 Initially, pits form randomly in the anode surface (unless promoted by stamping prior to anodization) due to the dissolution of alumina (Equation 5.6). Under the right temperature (kinetics and diffusion) and voltage conditions (electric field and migration), the alumina growth and dissolution eventually results in parallel pore growth into the aluminum work piece (82). This can progress until all of the aluminum is consumed, leaving a parallel array of pores open at one end and closed at the other.

After anodization, it is typical to expose the AAO to a stronger etchant, such as

H3PO4, for a brief period of time to widen the pores and open the closed end. Further details of the mechanisms of pore growth, diameter, spacing, and thickness of the alumina cap at the closed end of the pores are available in the literature (81; 82; 83).

We coated silicon wafers with a thin film (100 nm) of Pd, followed by a thick film (1 m) of Al, with adhesion layers of 5nm Cr or Ti between each deposition. By incorporating a conductive Pd layer, we allow the anodization to progress through the entirety of the Al layer, with no blocking alumina “cap” at the end of the pores. Additionally, the Pd film serves as the cathode for electroplating of Pd wires. We did not find evidence that the adhesion layer prevented anodization through to the Pd or plating after anodization. The Al layer was anodized in 3.75 wt% oxalic acid at 10°C with an anodization potential of 40 V against a lead sheet cathode. The bath was stirred continuously to help maintain a constant temperature. The anodization current was calculated from the voltage drop across a known resistor, which was recorded using a Gamry Instruments potentiostat. Figure 5.13 shows typical current-time behavior of the anodization. The shape of the curve in the figure is consistent with results in the literature for anodization (84). The small peak and sharp rise in current at the end of the chart is also seen in the literature and is attributed to etching of Si. Therefore it is possible that the adhesion and Pd layers are also thru-etched before this point in the curve.

Figure 5.14 shows a TEM micrograph of AAO prepared in the manner described above. The pore channels, visible as a light contrast stripes in the AAO layer, are approximately 30 nm in diameter with 100 nm spacing. It is not clear if the pores in this image reach the Pd layer; some of the channels on the right side of the image appear to stop short. However, we speculate that the semicircular regions of light contrast at the interface between Pd and AAO may open regions at the Pd interface with the pore channels that have been exposed by a slightly oblique ion beam milling angle.

Figure 5.13 – Current-time behavior of anodization. Sample bias 40V, 10°C, Pb cathode. Vertical scale: 60 mV = 10 mA. Sample area 1 cm2.

Figure 5.14 – TEM image of AAO on Pd supported by Si wafer. Image by Munekazu Motoyama.

5.4.2 Pd Electroplating in Laboratory-grown AAO

As previously stated, the primary motivation for growing AAO in the laboratory was to avoid having to build up a substrate on the AAO or attach the thin film supported nanowires to a substrate. Therefore, it is critical to the success of the laboratory-grown AAO that we are able to uniformly electroplate into the pore channels and control the extent to which the pore channels are filled.

The plating procedure for lab-grown AAO was the same as for the commercial AAO (sans the sputtering step, as the lab-grown AAO is formed on a conductive Pd layer specifically for the purpose of electroplating). In every case the liquid electrolyte plated into only a small fraction of the pores channels, resulting in much faster growth rates than anticipated in the channels that did plate, leading to caps forming on top of the template. Figure 5.15 shows typical results from attempts at Pd electrodeposition in lab-grown AAO. Efforts to improve the ratio of filled to unfilled pore channels included sonication, surfactants, and increased pore widening times.

Figure 5.15 – Pd deposition in lab-grown AAO pore channels. Pd deposits on the surface of lab-grown AAO, left, and a closer view after AAO has been removed by soaking in phosphoric acid, right. The hemispherical caps form when wires the pore channels in that region are completely filled. As seen in the image on the right, some regions exhibit Pd deposition as expected in the template, while others see no wire growth.

We have not yet demonstrated uniform filling of the pore channels, suggesting that significant variations in anodization rate on the sample surface resulted in regions of pore channels that did not reach the Pd layer. The presence of unanodized Al may be quite detrimental to the fabrication process due to the strong reactivity with HCl, the electrolyte used for Pd plating. Difficulty plating uniformly in laboratory AAO pore channels is not uncommon, and there are other reports of this issue in the literature (85). It may be possible to succeed with this method through meticulous cleaning of the sample before and after anodization, or with larger pores spaced further apart.

Figure 5.16 – Free standing AAO/BYZ membrane on a silicon wafer piece. AAO grown by anodization of Al with Pd thin film contact. BYZ deposited by pulsed laser deposition (PLD). It is not clear what the solid line running perpendicular to the AAO pores is. We speculate that

96 this sample experienced a change in deposition parameters during Al evaporation at this point. Image corresponds to * in process shown in Figure 5.17.

5.4.3 Fuel Cell Fabrication with AAO Grown on Substrate

We have demonstrated the feasibility of deposition of BYZ by PLD on the backside of the lab-grown AAO and observed the free-standing membrane of AAO/BYZ in the SEM, as shown in Figure 5.16. The fabrication process is based on one commonly used by our lab (66) with the addition of the steps necessary for AAO growth and Pd electrodeposition. The other noteworthy change is that the flat side of the silicon wafer becomes the anode rather than the KOH-etched side. The BYZ membrane is robust enough to support itself over the window area etched in the Si wafer, but the Pd film is far more likely to fail and therefore the additional reinforcement from the BYZ layer is needed before the AAO can be removed. We expect that this process is capable of producing thin-film fuel cells with nanowire electrodes should the problems of plating in the lab-grown AAO be resolved. The process for fuel cell fabrication with lab- grown AAO Pd wire templates is shown in Figure 5.17.

Figure 5.17 – Proposed fuel cell fabrication process for Pd nanowire anode SOFCs. Additional steps from standard procedures (66) used in our lab. When dissolving AAO, the BYZ layer must be isolated from the etchant solution or will also be removed. Figure 5.16 shows SEM of sample at * in process above, although Pd wires were not successfully plated.

5.5 Alternate strategies to high surface area MEIC electrodes

This section describes some strategies that do not require nanowire growth but may be promising for fuel cell fabrication or research purposes. Before we explored the use of nanowires for high surface area electrodes we briefly investigated photolithography methods. After encountering some frustration with electroplating in the laboratory- grown AAO, we investigated two strategies with the inverse structure of nanowires – first, ALD of Pd in AAO pore channels; and second, lyotropic phase electroplating. The processes, advantages, and difficulties of each of these methods are presented in this section.

5.5.1 Photoresist templating

Silicon wafers coated with 100 nm Pd by electron beam evaporation were primed with hexamethyldisilazane (HMDS). The primed wafers were then coated with 7 m of Shipley 220-7 positive photoresist. After exposure, the unmasked regions were removed by the development process. Wafer pieces were then electroplated with Pd from a 0.5wt% PdCl2 in 10wt% HCl solution and the photoresist was washed off with acetone. The result was an array of cylindrical features as shown in Figure 5.18.

Figure 5.18 – Pd pads electrodeposited through a photoresist shadow mask. Nonuniformity of pad height distribution can be observed in the above SEM. Rough appearance of top surface is typical for electroplated the Pd from PdCl2 in HCl in our lab. A closer look at the side walls shows relatively dense Pd deposition.

Based on the spacing, diameter, and maximum height of the cylinders, this sample has a roughness factor of approximately 2, which is expected to have a small impact (less than sample-to-sample variation observed in our lab) on fuel cell performance. It is not clear why some of the exposed areas experienced significantly lower deposition rates. One possible explanation is incomplete removal of exposed photoresist during development. If a sufficiently thick photoresist is used, this strategy may work to enhance the anode surface area by an order of magnitude. However, photolithography is typically limited to feature sizes on the order of microns, and thick resist tends to have a larger minimum feature size, thereby limiting the maximum roughness attainable. The 5 m features in Figure 5.18 are therefore the practical limit for the photoresist-masked Pd pillars.

5.5.2 Atomic Layer Deposition of Pd nanotubes

The AAO templates used are brittle, but, at 50 m, strong enough to handle if a fuel cell was built on one side. This approach uses ALD to deposit a thin film of Pd on the walls of the template pores, as reported by Elam et al. (86). Pd ALD is difficult, and we were unable to uniformly coat the pores of an AAO template with the ALD arrangement used. The device structure is shown in Figure 5.19, along with I-V and power density results showing a peak power density of roughly 3 W/cm2 at 400°C. A Pt catalyst layer of ~ 2nm was also deposited by ALD inside the pores prior to deposition of the Pd film.

The potential surface area increase of the ALD-coated AAO is nearly equivalent to that from nanowires grown using the AAO as a template. This is because the side-wall area is estimated to be the same, but only a portion of the initial substrate area is preserved as the AAO covers the space between pore channels. However, we expect that this device may suffer from higher impedance to H conduction in the Pd as the cross-sectional area of Pd available for diffusion is significantly smaller (pore circumference times Pd thickness compared to pore channel cross-section). The power density of the device is insignificant compared to the thin film planar structures studied in our lab (66), but the fact that a fuel cell I-V behavior was obtained (Figure

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5.19) suggests that with corrections to the ALD process such as increased exposure time between precursor pulsing and chamber pump/purge to allow the precursor flow through the entire length of the pore channels, this device may show improvement. It is not clear why the OCV is low (only ~ 0.75 V), but this may be due to improper sealing between the device and the alumina tube used to isolate the anode from the cathode environment.

Figure 5.19 – ALD Pd deposited in AAO pore channels for SOFC. Schematic diagram (top) of AAO with Pd deposited by ALD in the pores, and the architecture of a fuel cell based on ALD-coated AAO. SEM image of the device (bottom left) and current-voltage and power density behavior (bottom right). The SEM shows brightest contrast near the top of the AAO pore channels, suggesting a shallow penetration depth of the ALD Pt catalyst.

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5.5.3 Lyotropic phase Pd electrodeposition

By mixing the amphiphillic molecule octaethylene glycol monohexadecyl ether

(C16EO8) with an aqueous solution of ammonium tetrachloropalladate, (NH4)2PdCl4, a hexagonal (H1) lyotropic phase is formed. The C16EO8 forms nanometer-scale columns that are hexagonally packed and surrounded by the aqueous phase. If the mixture (a very highly viscous solution) is placed between an electroplating cathode and counter electrode, the Pd can be electrodeposited onto the cathode surface, resulting in a dense Pd layer with parallel pore channels with nm-scale diameter. This process has been previously described by Bartlett et al. (87)

We have made preliminary progress in the formation of fuel cells with Pd anodes based on this process. Figure 5.20 below shows characteristic cyclic voltammograms from a clean Pd foil surface and a clean Pd foil electroplated from a lyotropic phase. The active surface area was approximately 39 times greater for the electroplated sample, based on the height of the surface oxide stripping peak. This is significantly lower than the area enhancement observed by Bartlett (87) for the same loading (~240x), most likely due to deviations in the proportions of components in the solution and post-deposition washing process. We expect improvement in future results with more experience with the procedure.

The samples were then each coated with a film of approximately 1 m thick BYZ by PLD. Porous Pt pads were deposited by DC sputtereing onto the BYZ surface for the cathode. The anode was not coated with any other materials, nor was it chemically treated to remove any oxide formed during PLD. Figure 5.21 shows the results of fuel cell testing. The maximum power density of the H1-phase electroplated sample was approximately 1.8 times that of the untreated Pd foil samples. It is possible that the fabrication process reduced the effective surface area of the H1-e Pd anode due to morphology changes at 600°C during PLD.

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Figure 5.20 – CVs of flat and lyotropic H1 phase electroplated Pd surfaces. Cyclic voltammograms performed in 0.5M H2SO4 at room temperature. The large positive peak represents the reduction of Pd surface oxide. The oxide formation is seen at high positive potentials and negative current. The inset shows the detail of the Pd foil curve. Scan rate 100mVs-1. The potential was not cycled to potentials negative enough to observe H adsorption and desorption peaks.

Figure 5.21 – Fuel cell behavior of SOFC with lyotropic H1 phase deposited Pd anode. I-V and power density behavior for flat Pd anode (blue) and electroplated nanoporous Pt anode (red). Data taken at 350°C.

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The results of this study are preliminary, and further testing is required to verify the results. Of particular interest are the effects of including a Pt catalyst and the impact of using methanol fuel rather than hydrogen. Significant volume changes of nano-sized Pd features are known to result from absorption of H and oxidation of Pd. Future studies of high surface area Pd structures should examine the effects of heat cycling and fuel cell testing on Pd morphology.

5.6 Summary

We have studied various means of producing high surface area electrodes from mixed conducting materials for thin-film solid oxide electrolyte fuel cells. Many of the paths explored led to significant difficulties in the fabrication process. However, with some further study of experimental conditions, we suggest that many of the device fabrication strategies discussed above can be successful. In particular, plating thick layers of material onto commercial AAO to build a substrate and lyotropic phase electrodeposition onto foil or silicon-based substrates appear to be promising.

It is important to reiterate that the techniques discussed in this chapter are not limited to palladium electrodeposition, but could potentially be used for other electrodeposition-friendly materials, such as RuO2, and therefore allow for the development of high surface area cathodes for low- and intermediate- temperature SOFCs, potentially lowering activation losses for both anode and cathode processes.

104

6 Conclusions

Fuel cells present an opportunity to meet energy demand in commercial, industrial, residential, transportation, and military sectors as a replacement for fossil fuel based energy as well as battery storage in some scenarios. Among the principle disadvantages to using fuel cells are the disproportionately high cost per watt relative to combustion and battery technologies, as well as the very high operating temperatures typical of some electrolyte materials. The goal of much of the work done with fuel cell technologies in our lab is to lower the cost and temperature of fuel cells by increasing current and power density characteristics of the various fuel cell systems studied. This work presents efforts to reduce activation losses in fuel cells through surface modification of metal electrodes (§3 and 5), as well as study the behavior of fuel cells under new operational conditions (§4).

For the “bioelectricity” device discussed in §3, we examined the kinetic response to electrode surface treatments. By attaching a monolayer of a molecule to an electrode, the interaction between the charge carrier ferredoxin (Fd) and the electrode is changed. When that interaction is sufficiently similar to the reaction mechanism that reduced Fd undergoes in nature, the charge transfer rate (i.e. the probability of charge transfer) increases dramatically. The effects of the surface chemistry on the Fd oxidation reaction rate can then be extrapolated to determine the expected behavior for a fuel cell-like device through activation loss estimates. In particular, we found that a poly- L-lysine coating resulted in nearly an order of magnitude increase in the heterogeneous rate constant for electron transfer from reduced Fd to a gold electrode, resulting in significantly decreased activation losses, and higher predicted power output for a hypothetical device. Recent progress by others in our lab suggests that such a device is possible, but the predictions regarding light absorption and conversion to electricity show that the device would not be competitive with present-day commercial silicon solar technologies. The underperformance relative to commercial PV and the complicated, probe-at-a-time fabrication process, strongly suggest that for

105 the bioelectricity arrangement discussed to become a competitive energy technology, significant design changes must be considered.

In §4, the results of a study on a direct methanol fuel cell (DMFC) were presented. This study is interesting for the use of a ceramic oxide-conducting electrolyte at relatively low temperatures for a solid oxide-based fuel cell. Furthermore, the methanol was delivered without reformant. We found that the fuel cell behavior suffered slightly from the use of methanol instead of hydrogen fuel, and that a significant portion of the fuel cell impedance is likely attributable to the use of a PtRu co-sputtered alloy anode. Despite this impedance, the overall performance of the devices in H2 fuel was comparable to previous results from our lab using similar fuel cell architectures. A PtRu alloy is often used in DMFCs to alleviate the problems of Pt-CO “poisoning” through reaction with Ru-OH in the “bifunctional mechanism”. It is not clear what reactions occur on the anode side, and whether the bifunctional mechanism reaction route is possible in the absence of reformant delivered with methanol, or if the Ru-OH species forms from O2- diffusion through the electrolyte. By developing an understanding of the fuel cell behavior in this system, we hope to be able to improve performance by selectively targeting sources of loss.

In the final study presented, §5, the fabrication of solid oxide fuel cells (SOFCs) with high-surface area anodes is discussed. We elected to explore the use of mixed conducting materials that can be deposited through electroplating. By depositing material such as Pd through a nanostructured template, a fuel cell electrode with high surface roughness can be fabricated. Although we encountered numerous process difficulties in fabrication, we were able to demonstrate that the anode surface area can be increased sufficiently to improve fuel cell performance. While this is somewhat counterintuitive for H2/O2 fuel cells where the ORR tends to limit performance, the results are promising for the case of methanol oxidation on the anode, which, as seen in §4, may contribute significantly to loss. With proper materials selection and process design, many of the strategies discussed for increasing reaction surface area are also applicable to the cathode side, and it may be possible to very significantly reduce

106 activation losses in low-temperature ceramic fuel cells. This would in turn decrease the cost per watt and perhaps allow for lower-cost (non-PGM) catalysts.

107

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