A life in electrochemistry.

Annual Review of Analytical Chemistry 2014.7. Downloaded from www.annualreviews.org by University of Saskatchewan on 05/27/14. For personal use only.

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A Life in Electrochemistry Allen J. Bard Annual Review of Analytical Chemistry 2014.7. Downloaded from www.annualreviews.org by University of Saskatchewan on 05/27/14. For personal use only.

Department of Chemistry, University of Texas, Austin, Texas 78712; email: [email protected]

Annu. Rev. Anal. Chem. 2014. 7:12.1–12.21

Keywords

The Annual Review of Analytical Chemistry is online at anchem.annualreviews.org

biography, research, electroanalytical chemistry, chemistry

This article’s doi: 10.1146/annurev-anchem-071213-020227

Foreword

c 2014 by Annual Reviews. Copyright All rights reserved

To tread a newly opened furrow, to feel complete freedom of action, and to see on all sides new vistas opening for research, brings with it a joy that only those who have had the raw pleasure of original research are fully in a position to grasp. Henri Moissan (1)

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Work in electrochemistry over more than two centuries has had a truly revolutionary, and sometimes underappreciated, role in science and technology. The first period (Volta period) introduced the first reliable source of electricity in an area that remains important in searches for new batteries for devices, propulsion, and grid storage. The second (Davy period) showed how many new elements and compounds, not previously available in chemistry, could be prepared using electricity (e.g., fluorine). The third (Nernst period) connected electrochemical measurements to thermodynamics and solution physical chemistry. The fourth (Heyrovsky period) introduced unique electroanalytical methods, especially for metal ions. The most recent period, which dates from approximately the middle of the twentieth century, connected the field to modern organic, inorganic, physical, and biochemistry and to theoretical concepts such as molecular orbital (MO) theory and mechanisms of outer and inner sphere electron-transfer reactions. I appreciate the invitation of the editorial board of the Annual Review of Analytical Chemistry to prepare an autobiographical article. After some brief personal details, I try to chronicle this fifth period of electrochemistry, especially how our group contributed to its development, and also perhaps to offer some lessons I’ve learned in the process.


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EARLY LIFE Although this is mainly an autobiography about my career in science, I’d like to include a little about my personal life, because I think it had a critical role in shaping my work and me. I was born December 18, 1933, in New York City in the midst of the Great Depression that began approximately four years earlier and only really ended with the start of World War II approximately eight years later. My family had only recently moved to New York from Minnesota, where my brother and sister, 11 and 10 years older than me, grew up. On September 14, 1906, at age 14, my father came to the United States by himself from Gomel, Belarus, where his father, Raphael Mordechai Barishansky, was the chief rabbi. After arriving on Ellis Island in New York, he immediately went to Minnesota where he had a distant relative. My mother came years later with her family from Latvia. My early years were spent living in an apartment in the Bronx across the street from Crotona Park, which was a great place for a child to explore. That, and the nearby Bronx Zoo and Botanical Gardens, came as close to the “outdoors” as was possible in a big city. In a sense, I had the best of both worlds, because I could also take the subway downtown and visit the museums, which I did very frequently. My sister and brother played an important role in my early education. My brother had big chemistry and erector sets, as well as other toys and games that I could get into, even at a young age. My public school career (K–9th grade) was pretty unremarkable. My teachers, by and large, were very good and cared about the largely diverse group of kids in my neighborhood. Science was my favorite subject, and I loved to look up articles about science (mostly about plants and animals and also a little about other things that I could understand) in the encyclopedia The Book of Knowledge (2). I, as well as almost all my friends, also liked to build things and so, in these pre-TV, pre-computer (“pre-everything,” as my grandson tells me) days, we spent lots of time building things and carrying out other projects, such as collecting and identifying plant or insect specimens. After World War II started, my brother went into the army and was stationed for a time in Fort Hood, Texas, and he would send me leaves from trees in Texas to add to my collection. I also liked to try my own experiments, most of which didn’t turn out the way I wanted. I remember that when I walked to school, particularly on rainy days, there would be multicolored oil spots in the streets where cars parked. I liked to paint with watercolors but had heard about oil paint and had the idea that somehow those spots were oil colors. So one day I took a paintbrush and piece of white cardboard and tried to pick up the colors in the oil slick. I was disappointed when 12.2

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I couldn’t transfer any of the colors to the cardboard, and what I painted was basically colorless. I asked why this could be but didn’t get any satisfactory answers; I did understand, at least, that there were ways of getting visible colors that were unrelated to the presence of dyes or pigments. A second unsuccessful experiment I remember was my attempt to devise a perpetual motion device. This started because I had become interested in tropical fish and had managed to get my father to support this hobby as well. As was my usual style, I read everything I could find on the subject, including a rather sizeable monograph called Exotic Aquarium Fishes (3), and then my father and I proceeded to put together as many aquariums, typically 15-gallon ones, as we could manage. This was a good experience, because it meant learning about pumps, filters, balanced biological systems, not to mention the sexual life of fish. I got pretty good at setting up rather eye-pleasing aquariums and would go to the stores in lower Manhattan along Nassau Street to see what inexpensive plants and fish I could afford. The need to drain these large, heavy aquariums necessitated learning to use a rubber tube siphon. In addition, I liked to experiment with different setups. I would siphon the contents of our sink into a gallon bottle on the floor to see how long it would take to fill the bottle. Then I got the idea that I could make the water flow uphill at the same time by putting a two-hole stopper into the bottle with glass tubes through it, one going all the way to the bottom of the bottle and one ending right below the stopper. If I connected the shorter glass tube to the siphon from the sink and the longer glass tube to a rubber hose leading back up, as water siphoned into the closed bottle, the increasing air pressure above the water filling the bottom of the bottle would force that water up into the tube leading back to the sink. The idea was to have the system continue to flow without any external energy. Of course, it didn’t work. I could get the water to flow up almost to the level of the water in the sink but not quite far enough. However, these and many other experiments taught me a useful lesson (see sidebar, Experiments That “Don’t Work”). While in junior high, I became interested in photography and developing black and white film myself. It was easy to buy a few chemicals at a photography shop (developer, fixer) and with three trays and a red lightbulb for orthochromatic film, develop a roll of film to get negatives and then print them on photographic paper. I bought a reasonably good used camera and a tripod, and I would take pictures at the Museum of Natural History and develop them in the bathroom. I learned a little about the chemistry involved but really didn’t understand it very well, especially with regard to the latent image formed during exposure of the film and sensitization of silver halide. When it was time to go to high school, I was determined to go to one of the well-recognized, high-quality, and specialized schools in New York. The obvious one was the Bronx High School of Science, which was the hardest one to get into. Fortunately, I passed the exam and for the next three years would travel, first by trolley and later by bus, up to the old building near the Grand Concourse to be among the best teachers and students I had ever seen. It was at Bronx Science that I really gained a love and respect for chemistry, along with math and physics, and was able to take two years of chemistry courses and even get a volunteer job helping with the chemistry storeroom. While I was in high school, I got a job at the Physician’s Clinical Laboratory, which was mainly devoted to doing urine analysis for glucose or pregnancy tests. My job first involved

EXPERIMENTS THAT “DON’T WORK” You can learn a lot from experiments that don’t turn out the way you expect (that “don’t work”), often more than from ones that do.

www.annualreviews.org • A Life in Electrochemistry


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going around the Bronx by bicycle or bus and picking up samples left at the different pharmacies. However, I also liked to watch as the owner of the lab did the testing, and he eventually let me do the glucose tests. This was pretty simple chemistry: Add Benedict’s solution (copper sulfate with sodium carbonate and citrate) to urine in a test tube, heat, and see if the blue color changed to orange with the reduction of Cu2+ to copper oxide. He also did pregnancy tests by injecting urine into rats and then dissecting them. Quite a difference from the simple test that is now available.

COLLEGE DAYS


Approaching graduation from Bronx Science ( January 1951), I decided to apply to CCNY (formally the City College of the College of the City of New York). New York City had several colleges that were all pretty good and didn’t charge tuition. One had to take an examination to get into them, but by living at home and commuting to college every day, it was possible to get a good education. New York state didn’t have a state university system at the time (but later introduced the SUNY system), but they did give four-year state scholarships to deserving students, and that helped me buy books and pay lab fees; with part time work and summer jobs, I could get through college without amassing the obscene debts that modern students acquire. It is always amazing to me how the society in those times could afford this system, which made an amazing difference in my life (and my sister’s and brother’s years earlier) and that of many of my classmates as well, whereas our current affluent society cannot. I was very happy at CCNY and got an excellent education there. Of course, I enrolled as a chemistry major and was surrounded by a bunch of intelligent and motivated students and excellent teachers. Because CCNY had no graduate program, the professors taught all of the classes, even the labs, and were always available to the students. I can’t name all of those who deserve recognition, but I particularly remember Barnet Naiman, who gave me a real appreciation of analytical chemistry, and Walter Miller, who introduced me to physical chemistry and instrumental analysis. Up to then, I thought chemistry was just carrying out reactions, even to perform analysis, but I started understanding how new instrumentation can be used to characterize chemicals and study their behavior. There weren’t many instruments available at CCNY at the time, but I was introduced to polarography and the dropping mercury electrode and was fascinated by the technique and the ability to do electrochemistry and learn about a solution this way. In addition to chemistry, I learned a lot of other things in college, including an appreciation of the arts and humanities. It was great living in New York City as a college student. We might not have had a big green campus or a football team, but the city had affordable music (e.g., at Lewisohn stadium right on the CCNY campus) and theater. I always worked during the summers and was fortunate in 1954 to get a job as a student aide trainee first at the Naval Research Laboratory (NRL) in Washington, DC, and then at the New York Naval Shipyard in Brooklyn. The one at NRL was more in metallurgy and I was really more interested in chemistry, because I was eager to learn as much as I could about techniques and instrumentation. My job at the Brooklyn Navy Yard involved several projects. The one I remember best was gas analysis on aircraft carriers. Fires on the deck elevators that moved planes from one level to another had ignited as a result of vapors in the hydraulic systems reacting with oxygen in the systems. The Navy decided to replace the air in these systems with nitrogen and brought the carriers into the shipyard to flush the oxygen out of the elevator hydraulics. It was necessary to know when the oxygen level was low enough, and so one had to analyze daily the exit gas stream for oxygen. This was harder than it sounds. The bottom of the deck elevators were in the bowels of the ship, and the best available method of measuring oxygen gas was with the Orsat apparatus. This was a large, bulky glass apparatus with different glass chambers. One chamber 12.4

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EXPERIMENTAL APPROACHES


The most obvious approach to an experiment may not be the best way to do it. It pays to consider alternatives.

held the gas sample, and the others held liquids that would react with components of the gas; for oxygen, it was the pyrogallol solution. One took a sample of gas of known volume then flushed the gas back and forth into the pyrogallol solution where oxygen reacted; thus, the gas volume (at the same temperature and pressure) was now smaller. The measurement was pretty straightforward, but maneuvering down the narrow stairways carrying the fragile apparatus housed in a big wooden box was pretty difficult. After a few daily trips, we started thinking about an easier way, such as measuring the thermal conductivity of the gas mixture instead. Thermal conductivity detectors were beginning to be used more frequently in gas chromatography, where the resistance of a hot wire, which is a function of its temperature, is measured. Because the temperature is a function of the rate of heating and the rate at which heat is carried away by the gas, it is a function of the composition of the gas, and indeed, if one calibrated it with nitrogen-oxygen mixtures, one could use this technique to find the amount of oxygen. The needed apparatus to do that measurement was much lighter and smaller, and the measurement could be made faster, so it made things much easier (see sidebar, Experimental Approaches). I got my BS in Chemistry in January 1955 and immediately started thinking about graduate school. However, because that wouldn’t start until September, I wanted to find a job, preferably in chemistry, as soon as possible. A few friends and I took a road trip across the country heading south and then west. It was my first taste of the western United States and the first time visiting Texas and California. As soon as I got back to New York, I applied for a job at the General Chemical Company (an affiliate of Allied Chemical) in Morristown, New Jersey. Fortunately, I got the job and was placed in the analytical chemistry laboratory. That was great for learning new things, and it was the best-equipped lab I had ever seen. There were several fellows my age from New York City, and we could carpool daily. We all worked on different projects that arose in the research labs of the company, so one could learn a lot about different analytical methods and techniques, e.g., Parr bombs for digesting samples and atomic spectroscopy (still using photographic recordings of spectra). General Chemical made fluorocarbons that were competitors of DuPont’s Freons, and one project was to see if they could be used in aerosol cans with food products. A product of interest was pancake batter and the analytical problem was to see how much fluoride ended up in the pancakes- a pretty challenging problem, especially given that fluoride-specific ion electrodes had not yet been developed.

GRADUATE SCHOOL DAYS I had a pretty good record at CCNY and didn’t think I would have much difficulty getting into a good graduate school. The problem was how to afford it. In 1955, the now universal culture of supporting STEM (science, technology, engineering, and mathematics) graduate students and paying their tuition from research grants was not yet in place, and chemistry graduate students were treated the same as those in English and philosophy. The two best prospects for me were Columbia University in New York City and Harvard University in Cambridge, Massachusetts. Columbia was an excellent school and lots of CCNY graduates went there, but I thought I needed to leave New York and learn about another environment, such as New England. Earlier students from the City Colleges had done well at Harvard, so I was lucky to get a teaching assistantship www.annualreviews.org • A Life in Electrochemistry


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and a small scholarship, which would make it possible to live in a dorm and pay tuition, at least for the first year. So in September 1955, I arrived at Harvard and was determined to learn as much as I possibly could from the very distinguished faculty and my fellow students. One thing I learned quickly was that, although I had a good education at CCNY, it hadn’t covered a lot of the latest work in science. I knew no quantum mechanics and nothing about molecular orbitals, for example, whereas my fellow students from MIT and Columbia were all familiar with these. A student had to take qualifying exams in the first semester and also select a research area to do lab work in. This would often end up being the dissertation theme. I knew I didn’t want to be in organic or biochemistry, so I studied what was being done in the other areas. I was really interested in the recently discovered compound ferrocene and the possibilities of a new field of organometallic chemistry being pursued by an assistant professor of inorganic chemistry, Geoffrey Wilkinson, so I signed up to be in this lab. I was assigned to make CpCr(CO)3 Me with lots of help and advice from an older graduate student in the lab, T.S. Piper. What I hadn’t even considered was that Wilkinson was in a nontenured position near the end of his assistant professor term in a department that only rarely granted tenure to junior faculty. However, Wilkinson had done so well that it seemed that getting promoted with tenure was a good possibility for him. It didn’t happen, however, and Wilkinson left Harvard for Imperial College in early 1956, as I recall, and I had to find another lab and topic for my studies. Truth be told, I wasn’t so taken with the organometallic chemistry, even with the exciting developments in ligand field theory and new systems. I didn’t like doing synthesis that much, especially working in a glove box. The possibilities for me were thus either in physical chemistry, where Harvard had some really distinguished faculty such as George Kistiakowski and E. Bright Wilson (as well as some brilliant younger people such as Bill Klemperer and William Moffitt). However, I really liked analytical instrumentation, and James J. Lingane was a recognized leader in the field of electroanalytical chemistry. JJ was a former student of Isaac M. Kolthoff, who was probably the leading analytical chemist of his time in the United States, and many outstanding faculty in US universities, such as Herb Laitninen at the University of Illinois, who was JJ’s contemporary at the University of Minnesota, had worked with him. The lab at Harvard was located in a separate small building, Coolidge Hall, where I believe Theodore W. Richards, the first American Nobel Prize laureate (1914), had worked, and I learned in the course of my stay there that some of his chemicals and pieces of apparatus were stored in the stock room in the basement. JJ’s office and lab were on the first floor, in addition to a small conference room. The student labs were on the second floor, and when I joined the group, the more senior graduate students were Fred Anson, Don Davis, and John Kennedy, all of whom went on to have academic careers. I had to get up to speed in the new field, and I did this by reading Lingane’s (4) first edition of Electroanalytical Chemistry and also Kolthoff & Lingane’s (5) famous monograph Polarography. I also got lots of help and advice from my fellow students, especially Fred. JJ was mainly interested in coulometric titrations at this time, as well as voltammetry of metal ions, and I began work on the coulometry of tin ions in a bromide medium. Coulometric methods were interesting, because they were absolute analytical methods that didn’t depend on any calibration curves and yielded highly accurate and precise results (∼1 ppt). Unfortunately, they never really caught on as general analytical methods, although there were some important niche applications, such as determination of water by the Karl Fischer method. However, as I learned more and more about modern electrochemistry, it became clear to me and my colleagues in JJ’s group that the field was changing, and we wanted to be part of the change. The invention of polarography by Heyrovsky in 1924 and the use of the dropping mercury electrode introduced a new and important method in electrochemistry. However, by


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1956 the technique was reaching maturity, with its main impact in the analysis of metal ions and extraction of thermodynamic information for systems such as coordination compounds. The impact of the application of chemical thermodynamics in electrochemistry and analytical problems in general, driven by scientists such as Kolthoff and Furman at Princeton, was now well established. Modern trends in chemistry, e.g., quantum chemistry and MO theory, electron-transfer theory, physical organic chemistry, organometallic chemistry, and reaction mechanism studies, were yet to be applied to electrochemistry. Moreover, we were on the threshold of technologies such as the computer (Harvard had a “big” computer, probably less powerful than the Apple 2+, which was still 20 years away) housed in a building right in back of my dorm. This was still the vacuum tube era, but solid state electronics were starting to be developed. Most importantly, scientists such as Paul Delahay in the United States and Randalls in the United Kingdom showed how new instrumentation and methods such as chronopotentiometry or linear sweep voltammetry could address questions at shorter times than was possible by polarography. Delahay’s (6) book New Instrumental Methods in Electrochemistry became my new source of inspiration. I did have the more immediate problem in the spring of 1956 of supporting myself for the next years. I anticipated getting by on a teaching assistantship and perhaps working during the summers, but fortunately I got a National Science Foundation (NSF) student fellowship that would support me through the next two years. In 1957, Dave Geske arrived as an assistant professor at Harvard working with JJ in the teaching of analytical chemistry. He moved into the lab next to mine in Coolidge, and we had a wonderful interaction during my last year at Harvard. Dave did his graduate work at Iowa with Alexander Popov and was studying electrochemistry in acetonitrile (MeCN). (MeCN as a medium for voltammetry was also being studied at the same time by Kolthoff and Coetzee in Minnesota.) I was very impressed with this work, and discussions with Dave made it clear that use of an aprotic solvent that didn’t involve proton transfers coupled to the electron transfers could simplify the study of organic reactions by electrochemical means. At the same time, I had become interested in ways to study electrode reaction mechanisms by various electrochemical techniques. Delahay and his group had made a lot of progress, at least for first-order reactions with chronopotentiometry, and I thought that coulometry, a technique I knew well, could also make a contribution. Although it was fundamentally a slower technique, it was easier to obtain precise experimental results, and the mathematics of treating complex reaction mechanisms was easier. So while I was working on my dissertation research, Dave and I started working on studying reactions by controlled potential coulometry, and Dave was obtaining experimental results in MeCN to treat using this technique (7). In a way, my last year of graduate school working with Dave was like a postdoctoral position but in parallel with my graduate work.

OFF TO TEXAS I received my PhD in June, 1958, and in the previous fall I started looking for an academic position. Although I had other possibilities, I decided to accept a position at The University of Texas (UT) in Austin. Why would a New Yorker go to Texas? I wanted to see what “the West” was like, given that I had this romantic concept of the wide open spaces and a very different culture than I grew up in. I also thought that the times were changing and that UT, then a rather mediocre place, had lots of potential, if the state would use its abundant natural resources to finance it. Moreover, the chairman of the chemistry department at UT, Norman Hackerman, who was also a well-known electrochemist, made me an offer by phone within a couple of weeks of my sending in my application. I asked if he didn’t want me to come to UT and give a seminar. He said “no” and that I should let him know pretty quickly whether I accepted the offer. I flattered myself, assuming this instant interest was a reflection of my very strong resume but later discovered that this was www.annualreviews.org • A Life in Electrochemistry


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THE PREVAILING MODEL MAY NOT BE RIGHT The prevailing wisdom is sometimes wrong, even if it is espoused by prominent scientists and is believed by most people. Don’t believe everything you read. If there isn’t firm evidence for an assertion or model, be cautious about becoming a member of a scientific sect or clique that reinforces one another’s beliefs and concepts.


the usual modus operandi at UT, which saved them paying travel expenses and letting candidates see the true state of the department! I started working on electrochemistry in nonaqueous solvents and on new electroanalytical techniques for probing reaction mechanisms. The hot technique initially was chronopotentiometry. [Measuring potential change over time with a constant current source (e.g., a high-voltage battery and a big resistor) and the theory behind it was pretty easy (8, 9).] We worked on a technique we named cyclic chronopotentiometry, where one cycles the current from one value to another and analyzes the response (10). Another emerging technique was cyclic voltammetry (CV), where one scans the potential between two limits and records the current response. The instrumentation and theory were both more difficult, and the information content was basically the same as chronopotentiometry, but the observed CV response made it easier to extract the information. I built an instrument based on running a potentiometer back and forth with a motor and recording the response on an x-y recorder, a newly available apparatus]. Within a short time, however, the theoretical problems of CV were solved by numerical methods by Jean-Michel Savéant, Irving Shain, and Richard Nicholson (and later by Steve Feldberg using digital computer simulation). Moreover, there was a growing appreciation of operational amplifiers as components of easily designed potentiostats and other electrochemical subsystems, led by Don DeFord and others. Armed with these techniques, one could study organic species in nonaqueous solvents such as MeCN and start learning about their reaction mechanisms (11).

ELECTROCHEMISTRY OF RADICAL IONS There were numerous academics active in this area. In the United States, the group included those mentioned above, along with Geske, Ralph Adams, William Reinmuth, Anson, Charles N. Reilley, Royce Murray, Don Smith, Ted Kuwana, and others. Gus Maki and Geske had recently shown that a small electrochemical cell could be placed inside the cavity of an electron spin resonance (ESR) spectrometer and signals from species with an unpaired spin, such as nitrobenzene radical anion, could be obtained and analyzed (in spite of being told it would never work, because one couldn’t pass a current into the ESR cavity and not mess up the signal). This actually turned out to be a hugely useful tool in analyzing electrochemical systems (see sidebar, The Prevailing Model May Not Be Right). At the time, the prevailing models in physical organic chemistry mostly involved two-electron (2e) processes with popular intermediates being carbanions or carbonium ions. Generally, the field resisted the idea of radical ion intermediates, even in the face of very clear electrochemical results and ESR (see sidebar, Undoing a Prevailing Model). There was more difficulty in generating radical cations (because they tend to react faster with nucleophilic impurities such as water. However, by using vacuum line techniques and glove boxes, the purities of MeCN and other solvents were sufficient to demonstrate stable radical cations, even for aromatic hydrocarbons, and the importance of one-electron (1e) transfers as a general mechanistic path, especially, but not uniquely, in electrochemistry was accepted. In fact, it often became necessary 12.8

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UNDOING A PREVAILING MODEL


It is difficult to undo a popular model. This takes time and may go through various stages, even (a) the idea is wrong; (b) it may hold, but only for some special cases; or (c) it’s obvious and anyway we thought of it first.

to invoke structural rearrangements or reactions of the 1e products to account for the overall 2e paths, and eventually this view was adopted by the organic chemistry community. The ensuing years saw a growing body of experimental results in line with the 1e model, for heterocycles, ketones, aromatic amines and halogen compounds and many others. And studies of these reactions led to appreciation of their role in dimerization and polymerization (12). These eventually led to work demonstrating electronically conducting polymers. An important conceptual advance at the time was the importance of MO theory in interpreting electrode reactions. For example, Al Matsen, a theoretician at UT, suggested a correlation between Huckel MO theory and ¨ the energy for the electron-transfer reactions of aromatic hydrocarbons, and this concept was developed by Michael Peover and many others. What emerged was a useful approach to correlating MO, electrochemistry, and ESR and insight into how changes in structure affected the responses (13, 14). At the same time, very significant theoretical treatments by Rudy Marcus, Noel Hush, and others of the rates of electron transfer between coordination compounds, and later more generally, introduced and explained, at least semiquantitatively, the effect of molecular structure and solvent on the rates of electron transfer. Things developed very rapidly, and an interesting consequence was that a group of electrochemists, led by Adams, proposed to the Gordon Research Conference (GRC) that a winter conference be held on the West Coast in 1964 (breaking with a long GRC tradition of only having conferences in the summer in New England) (Figure 1).

ELECTROGENERATED CHEMILUMINESCENCE Several factors came together at this time to make the rise of electrogenerated chemiluminescence (ECL) possible. The Army Research Office was interested in possible applications of chemiluminescence (CL) and started a program to fund research in this area. Although I’m not sure the work under this program ended up being useful to the military, it did ultimately result in CL light sticks and novelties, as well as important applications in clinical chemistry, as discussed below. Early work in this area, e.g., by Ed Chandross (who was my roommate at Harvard), demonstrated that redox processes could produce CL, and David Hercules showed the electrochemical generation of light with an aromatic hydrocarbon, although initially the mechanism wasn’t clear. We saw that this was probably a radical ion annihilation reaction, even though, at the time, many believed that the radical cations were not stable. Groups at Bell Labs and American Cyanamid became interested in the process, and for a few years, ECL was a hot topic; although, relatively few groups got into the field, mainly, I suppose, because it involved difficult experiments. However, fundamental research on ECL was fascinating and it provided a connection of electrochemistry with spectroscopy and photoluminescence (15, 16). By studying the energetics, one could see how to generate triplets of species, such as 9,10-diphenylanthracene (DPA), that were not accessible from the photogenerated singlet state (because the fluorescence efficiency was close to unity). One interesting aspect of this work was a study of magnetic field effects on ECL. Merrifield at DuPont had studied photoluminescence in solids and showed that a magnetic field would affect triplet-triplet annihilation, and we wanted to see if this effect also occurred in solution. Indeed, it did and it also affected the radical ion annihilation (a doublet-doublet reaction) (17). Thus, not www.annualreviews.org • A Life in Electrochemistry


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Figure 1 The first Gordon Research Conference on Electrochemistry (1964). The picture shows mainly speakers. ( first row, from left): Barry Brummer, James Hoare, W.R. Ruby, A.C. Riddiford, Allen Bard, William Reinmuth, Ralph Adams, Richard Buck, Manny Baizer, Paul Delahay, Phil Boddy, Manfred Breiter, Don Smith, Heinz Gerischer. Other selected attendees. (second row, from left) Ted Kuwana (3); (third row, from left) Robert Osteryoung (15), Isaac Trachenberg (16); ( fourth row, from left) Ernest Yeager (15), James McIntyre (16); ( fifth row, from left) Harry Mark (11), Stanley Bruckenstein (12); (sixth row, from left) Lucian Gierst (10), Philip Elving (12), David Geske (15). Also attending the meeting, but not in this picture, were Fred Anson, Henry Taube, Rudy Marcus, Leon Dorfman, and G.C. Barker.

only was ECL an aesthetically pleasing phenomenon, but it made an important connection to spectroscopy and photochemistry (Figure 2). However, when useful applications such as light-emitting cells (this work preceded solid state light-emitting diodes) or ECL lasers didn’t materialize, most of the groups gave up their research. Our group stayed in the field, now supported by the NSF, mainly because it was fun and we were determined to see if we could make it analytically useful in aqueous solutions. We had shown early on that ECL was highly sensitive and could be detected at very low concentrations in MeCN and similar solvents (18), but the market for analyses in such solvents was negligible. Aqueous ECL seemed improbable because all of the ECL active compounds, e.g., DPA, were insoluble in water, and the electrochemical window of water (the potential range where the solvent itself is not either oxidized or reduced) is small, so the needed radical anions and cations could not be generated in the same experiment. However, in 1966 we read a communication by Hercules and Fred Lytle about a CL reaction of Ru(bpy)3 2+ (bpy – 2,2 -bipyridine) upon addition of strong base to a solution of the 3+ species and this species looked like a possibility for ECL. Although the communication suggested that this wouldn’t work in MeCN, we decided to try it and indeed it produced very intense ECL upon annihilation of the +3 and +1 species (19). However, the reduction potential to produce the reduced form, Ru(bpy)3 + , was too negative for use in an aqueous solution where preferential proton reduction occurred, even in neutral solution, although the oxidized form,

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Figure 2 Electrogenerated chemiluminescence (ECL) photos. (a) Generation of rubrene ECL with alternating current excitation. (b) Generation of 9,10-diphenylanthracene ECL at a rotating ring-disk electrode, with the anion radical generated at the disk and the cation radical at the ring. The ring of light is brightest on the inside edge of the ring, where the annihilation reaction occurs.

Ru(bpy)3 3+ , could be formed before water oxidation. The problem of generating a strong reductant in water was finally solved approximately eight years later by adding oxalate, because we knew that this species produces the strong reductant CO2 − upon oxidation (see sidebar, Fundamental Science and Societal Impact). Thus, by oxidizing both Ru(bpy)3 2+ and oxalate in an aqueous solution, we predicted we would be able to generate ECL, and indeed, that was the case (20). Analytical ECL was very sensitive, because light was produced under conditions in which there was no scattered light or interference from fluorescent impurities (21). This system became the basis of a commercial system for clinical analysis, where Ru(bpy)3 2+ was used to label antibodies or other biomolecules with a sensitivity that rivaled radioactive tags, and numerous tests were developed, first by IGEN and later by Roche Diagnostics.

FUNDAMENTAL SCIENCE AND SOCIETAL IMPACT It takes a long time from conception and initial experiments to the development of practical applications (e.g., for ECL, almost 20 years). The expectation that one can realistically propose “societal benefits” in a proposal about new science is misguided.



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PHOTOELECTROCHEMISTRY


ECL provided a convenient connection between electrochemical reactions and emission of visible light. I started thinking about the possibility of the reverse process—light in and electricity and chemistry out. I had given a seminar on ECL at the University of Wisconsin, I think in the late 1960s, and during the question period, a member of the audience, who I later learned was Farrington Daniels, said something along the lines of it being easy to make light by inputting electricity, as in ECL; what was more important and more difficult is converting radiant energy to electricity. I thought about it, but didn’t see how this could work in an ECL system, where, even if an excited state could be converted to a radical ion pair, the recombination would be rapid and the efficiency would be low. However, Heinz Gerischer, as early as 1960, had shown some interesting properties of semiconductors as electrodes. I had obtained a piece of p-Si during a visit to Texas Instruments in November 1960 and tried some dark voltammetry with it but had difficulties because of surface oxidation. In1972, Kenichi Honda and Akira Fujishima suggested that illuminating single-crystal TiO2 in an electrochemical cell could produce a current and perhaps even split water to hydrogen and oxygen (although this was not demonstrated in this first paper but was later done by adding a solution bias with the cathode in acidic solution and the photoanode in a basic one.). This came at a time when the world was experiencing a gasoline shortage and realizing that our fossil resources were finite. This work triggered new research in the field of photoelectrochemistry (PEC) by several groups aimed at studying illuminated semiconductor electrodes and seeing if the interfacial electric field would separate the photogenerated electron-hole pairs with a reasonable efficiency and produce photocurrents and electrolysis products. These included groups such as Mark Wrighton’s at MIT, Art Nozik’s at Allied Chemicals, Barry Miller’s at Bell Labs, and several others. Again, I was fortunate to get into the PEC field early as it was just developing. One of the things we wanted to try were other oxide materials, and we thought iron oxide (hematite) would be an interesting possibility, given that iron was abundant and its oxide’s absorbance in the visible region was better than the UV absorbance of TiO2 and other large band gap materials (22). We also wanted to try out polycrystalline materials, which were easier to obtain, and we showed that they worked pretty well (23). These included powders, and we thought it might be possible to use these to oxidize organics and other species, simply by throwing TiO2 into the solution and putting it in the sunlight. This idea was met with a lot of skepticism, because one didn’t have the same electric field driving force at the small particles that one has in much larger single-crystal electrodes. However, it worked very well, and we showed that one could oxidize cyanide ion in solution in this way (24) and later carried out the PhotoKolbe reaction, where platinized TiO2 added to a carboxylate such as acetate would produce CO2 and a hydrocarbon under irradiation (25, 26). In fact, one could even do a Urey-Miller-type experiment and generate amino acids with illuminated platinized TiO2 in an atmosphere of methane, ammonia, and water (27). We also wanted to see how the semiconductors behaved in aprotic media, given that we had a lot of experience with these. The wide potential windows, which were so important in ECL, were also useful here, because one could map the energetics of a wide band gap semiconductor such as TiO2 and its n-type behavior (28). The model for this semiconductor, where electrons are present in the conduction band but not within the band gap, predicts that reduction is seen for solution couples with standard potentials within the band gap only at potentials at or more negative than the conduction band edge. However, we found reductions occurring for couples at potentials within the gap (more positive than the band edge), which we explained as the effect of the presence of surface states. This demonstrated that this approach is a good way to map the electronic energy levels of a semiconductor. We used this technique with numerous semiconductors, but it is not

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used much in current research in PEC, mainly because, I think, many of the practitioners are more interested in water splitting and aren’t very familiar with MeCN or other such solvents. In 1977, I was invited to be a Sherman Fairchild Scholar at Caltech. This was not only an opportunity to work with Anson and his group, but also with Mark Wrighton, then at MIT, who also had a visiting position at Caltech. Mark was also very active in PEC, and he and I had a lot of opportunities to discuss this work. I remember one conversation where we were discussing how one could understand the instability of some semiconductors such as cadmium sulfide (CdS) under irradiation in solution. We pretty quickly figured out from energy levels in the molecule and the energies of potential reactions such as oxidation of sulfide ion in CdS how one could predict these kinds of things. Mark came back in a short time with a paper on this that we submitted to the Journal of the Electrochemical Society. This turned out to be a fairly highly cited paper (29), perhaps a record for the number of citations per time spent on the “research.” Marye Anne Fox joined UT in 1976 and also became interested in semiconductor PEC. She independently showed interesting organic chemical reactions at illuminated semiconductors. We also collaborated in a large project supported by the Gas Research Institute that included Mike White, Tom Mallouk, and several other faculty members aimed at using semiconductors to produce hydrogen; this has become a very active research topic more recently, but at that time few were working in this area. One concept that emerged from the collaboration was the integrated chemical system. This is the idea that to carry out complicated processes one has to put together a system of different components, analogous to biological systems. For example, we put together a PEC system for using CdS as the light-capturing semiconductor with sulfide ion as the electron source to produce hydrogen (30). This was accomplished with a Nafion membrane, a polymer that contains exchangeable cations, as the support. Dipping this membrane into a solution of a cadmium salt put Cd2+ into the membrane, and then treatment with H2 S produced micron-sized particles of CdS, which turned the membrane bright yellow and introduced protons as the exchangeable cation. Dipping the membrane in methyl viologen (MV2+ ) exchanged this cation into the membrane. When this membrane was irradiated in a sulfide-containing solution, it quickly turned deep blue, because irradiation caused formation of MV+ with oxidation of sulfide. Although MV+ is reducing enough to produce hydrogen, the reaction is very slow. However, Pt is a good catalyst for this reaction, so introducing Pt into the membrane showed production of hydrogen (via oxidation of sulfide) (Figure 3). I used the theme of integrated chemical systems when I gave the Baker Lectures at Cornell in 1987.

ELECTROCHEMISTRY OF POLYMERS Another topic of interest in the early 1970s was modified electrodes. The idea was to examine monolayers of strongly adsorbed materials and design their composition and structure so that one could control and perhaps catalyze electrode reactions. Pioneering work in the field was carried out by Art Hubbard, Royce Murray, Fred Anson, and others. The electrochemical response from monolayers was clear but pretty small, and it was difficult to see that the relatively small number of adsorbed molecules could catalyze electrode reactions effectively to obtain a high current density. We had the idea of looking at the electrochemical behavior of polymers, both dissolved in solution and on the electrode surface. Inherent in this work with electroactive polymers such as poly(vinylferrocene) (PVF) was the question of what the voltammogram would look like when the polymer had numerous oxidizable (or reducible) groups. We had begun this work shortly before I went to Caltech on the Fairchild Scholar program, looking at poly(vinylnaphthalene) and poly(vinylanthracene) (31). At Caltech, we got interested in this and looked at the theory of the multielectron oxidation of PVF and showed that for noninteracting ferrocene groups on the chain, www.annualreviews.org • A Life in Electrochemistry


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a

b

c

d

1 µm

Figure 3 Experiments with the Nafion/cadmium sulfide (CdS) integrated chemical system. (a) Piece of Nafion polymer film (clear) and after incorporation of CdS ( yellow). (b) Scanning electron microscope image of Nafion-containing CdS. The scale marker is 1 μm. (c) The Nafion/CdS film immersed in a solution of methyl viologen and sulfide ion before irradiation. (d ) The same film after irradiation with visible light, showing photoreduction of the methyl viologen. When the same film contains Pt particles, the system in panel d shows hydrogen evolution under irradiation.

each oxidizes independently with the same microscopic Eo , such that the voltammogram looks like a 1e wave, even with the transfer of 76 electrons (32). One could also make films of the polymers on electrodes and study the rate of oxidation of the polymer itself or other soluble species (mediators) via the polymer and study the rates of charge transport through the polymer layers. Another interesting polymer at this time was Nafion, developed by Walter Grote at DuPont. This was a very stable (Teflon-like) ionically conducting polymer that was crucial in the development of fuel cells for NASA’s Apollo program. I was consulting for DuPont at the time and managed to get a little of a then commercially unavailable lower molecular weight Nafion that was soluble in alcohol. This was very useful in making films on electrode surfaces and then incorporating electroactive cations into them. One could work out the rate of charge transport through such films and even, if one incorporated Ru(bpy)3 2+ , produce light via ECL (33). This was probably one of the first examples of a polymer light-emitting electrochemical cell. One could also incorporate polymers 12.14

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in PEC cells; for example, poly(benzyl viologen) coated on p-silicon electrodes would evolve hydrogen under illumination (34). One could also use the Nafion film as a template for growing structures such as needles of oxidized tetrathiafulvalene (35). As such, in the early 1980s, we were busy with PEC, ECL, and polymer electrodes. However, at this time, the first reports of scanning tunneling microscopy (STM) appeared by Binnig, Rohrer, and coworkers at IBM. This looked like an intriguing approach to imaging electrode surfaces and perhaps also adsorbed molecules, so we decided to try this technique. This launched us into what later became known as nanoscience.


NANOSCIENCE, SCANNING TUNNELING MICROSCOPY, ATOMIC FORCE MICROSCOPY, AND THE SCANNING ELECTROCHEMICAL MICROSCOPE In 1985, commercial STM instruments were not yet available, so we set out to build one and then examine an electrochemical interface (36). What is amusing is that although the reported STMs were all based on analog instrumentation, we decided it was easier to build one based on an Apple IIe computer with D/A and A/D interfaces, probably the first reported digital STM, and also to try it on a Pt/solution interface. We couldn’t get atomic resolution in our early attempts but could image interdigitated Al structures approximately 0.8 μm high, with several different width and spacing patterns, deposited on an SiO2 layer on single-crystal Si and covered with a 100-nm sputtered Pt layer. Later versions of the STM could show atomic resolution, at least on C and Au surfaces, but after numerous studies, we became frustrated by the lack of real chemical information, and especially quantitative results that could be used to understand chemical reactivity. One could obtain impressive pictures, often with many trials and careful selection of results, but we didn’t think we were making progress in electrochemistry. Since our earliest STM experiments (36), we had been thinking about a related scanning probe technique and began developing the scanning electrochemical microscope (SECM) (37). This employed a small electrode tip (a so-called ultramicroelectrode, or UME) with dimensions in the micron regime that was scanned over a sample, either a conductive material or an insulator, in an electrolyte while recording the electrochemical response. The background of UMEs, e.g., the work of Fleischmann and Mark Wightman, and the development of the SECM is described elsewhere (38). SECM has turned out to be very useful, not only in obtaining topographic images, as the STM, but also in measuring the kinetics of fast homogeneous or heterogeneous reactions and in many other studies of polymer films, corroding materials, liquid/liquid interfaces, and electrode surface structures. It could also be used to modify surfaces by electrodeposition and etching. A particularly interesting aspect of SECM is positive feedback (i.e., amplification) of the tip current when it is above a conductor. Generally, the current at a UME when it is far (several tip diameters) away from any surface is a steady-state current representing the flux of the electroactive reactant, O, to the tip and the flux of product, R, away from it. However, when the tip is close to a conductive substrate, R is converted back to O, which then diffuses back to the tip, giving a current that is larger than that given when it is far from the substrate. In fact, this feedback effect increases as the distance between tip and substrate decreases. Moreover, if that distance is reduced to the order of ∼15 nm, the amplification is so large that even a single molecule can be detected as it cycles around in the tip-substrate gap, and one can see effects of the molecule as the electroactive molecules diffuse into and out of a constrained gap space (39) (Figure 4). Although single molecules can be studied by spectroscopic, e.g., fluorescence, techniques (W.E. Moerner and my colleague Paul Barbara), electrochemistry offers a complementary technique to spectroscopy by probing different molecular properties. Barbara and I later combined the two techniques to study single molecules and polymer particles and showed that the electrochemical www.annualreviews.org • A Life in Electrochemistry


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a

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b 15 nm

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e

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Time (s) Figure 4 Single-molecule trapping and electrochemistry experiment. (a) Schematic diagram of the electrode (the radius in panel a) trapping species A between tip and substrate spaced a distance d. A is oxidized at the tip and reduced at the substrate. (b) Curve 1: Time history of current with tip (a ∼ = 7 nm) separated by d ∼ = 10 nm from indium tin oxide substrate in solution of 2 mM of Cp2 FeTMA+ (species A) and 2.0 M of NaNO3 . Curve 2: Response in 2.0 M of NaNO3 alone with the tip within tunneling range. From Reference 42, copyright 1996 by the American Chemical Society.

response for an immobilized species could also be recognized by the quenching of fluorescence upon a redox process (40, 41). We also used the SECM to examine ECL and PEC systems. For example, it was useful to prepare arrays of PEC candidate semiconductor materials of different compositions with a robotic dispenser and then use SECM with the tip replaced by a fiber optic light source to scan the array and examine the relative photocurrents generated at each spot (43). This not only allows a rapid method of finding a good material, but also allows one to see how various dopants affect the efficiency of the semiconductor for generating products. A similar approach can be used for detecting electrocatalysts, e.g., for the oxygen reduction reaction (ORR), which is an important reaction in fuel cells and certain batteries. Using this approach, we were able to find a Pd-Co-based catalyst for the ORR in acidic solution that was comparable to that of Pt, the best-known electrocatalyst. We also continued work with the STM and atomic force microscopy (AFM). For example, with STM, we could look at the formation and evolution of etch pits during heating of highly oriented pyrolytic graphite in air, and could image where etch pits grew on the basal plane at locations where atomic defects or catalyst particles existed (44). AFM could be made more quantitative, and we employed it to map the forces at the electrode/solution diffuse double layer (45). We also became more interested in metal, semiconductor, organic, and dielectric NPs in ECL and PEC and as described later, in single nanoparticle electrochemical experiments.

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we also explored other interesting scientific problems. A very cursory overview is given here, especially about a few papers that didn’t fare as well as I think they deserved. We did a fair amount on examining the behavior of thin films of a liquid crystal zinc porphyrin (ZnPor), in another joint project with M.A. Fox. The idea was to see if one could synthesize a well-ordered material by attaching long chains to a porphyrin and whether this would show a photoresponse as a single film sandwiched between transparent conductors (46). This was produced by heating the material to the liquid state, introducing it between the electrodes, and allowing it to cool to the solid single crystal. There was a belief that a single homogeneous film without any asymmetry would not show photogenerated electron-hole separation. In fact, it did, because of differential rates at the interfaces, but the efficiency was not good (47). We later showed that if one applied an electrical field to the film during the cooling stage, the photocurrent was much improved, at least initially (48). We also suggested how the film could be purified by zone melting of the film itself. Later experiments showed that this same film could be used by applying a potential and light to trap charge in the film in very small volumes, with the guess that one might be able to isolate the charge in a single ZnPor stack and thus have memory storage (49). However, we could only address the film with an STM tip, which wasn’t practical for random access memory. We also did a fair amount on electrochemistry in low temperature solvents, like NH3 and SO2 (50, 51). We also examined others, like MeCN and water, at high temperatures and pressures as supercritical fluids (52). NH3 has a very large window into the negative potential region, so one can study some interesting reduction reactions. For example a gold cathode under reduction conditions can produce auride ion, Au− (that can be anodically deposited to form a Au film) (53). The low temperature of experiments in these solvents also slows down coupled chemical reactions and allows radical ions that are not stable in MeCN to be observed. SO2 allows one to look at the very positive potential region, limited by the electrolyte anion, such as SbF6 − . In this solvent, one can see the oxidation of the bpy units in Ru(bpy)3 2+ (54). Supercritical water (e.g., at 385o C and 300 bar) is very interesting but difficult to work with. With oxygen it becomes a terrifically strong oxidant, but it needs to be maintained under high pressure in a titanium tube, and getting a good reference electrode is difficult. Many of the advances in electrochemistry have come from using new solvents (and things like Li ion batteries depend on this), so even more exotic solvents will probably emerge. We also became interested in electrostatic charging of insulators (dielectrics). We got into this indirectly because we wanted to see if we could produce an electrochemical cell, e.g., a very small-volume cell, with only one faradaic electrode. That indeed was possible, by having the counter electrode be one blocked from the solution by an insulating film, so this electrode only charged as a capacitor, but didn’t pass any charge to the solution (55). However, we thought it would be interesting to charge an insulator by contact electrification [i.e., rub a piece of Teflon with poly(methyl methacrylate)] and see whether that could cause a redox reaction. Now the “accepted” mechanism of contact electrification of two dielectrics was that it was an ion (not electron) transfer effect. In fact, we were able to show in several different papers redox processes by this technique, so long as one applied very sensitive voltammetric measurements, but I’m not sure we convinced the community (56). We carried out numerous other electrochemical experiments that we thought were interesting but had little impact in changing accepted beliefs. For example, it was generally believed that Hg would “poison” all enzyme reactions [without making a distinction between Hg(0) and Hg(II)]. We made a Hg drop electrode, buried a thermistor in it so its temperature could be monitored, adsorbed urease on the Hg surface, and showed that when it was exposed to urea, the temperature rose and the catalytic reaction occurred (57). We also showed that if the potential was scanned negative (to reduce disulfide bonds in the protein) the reaction stopped, but if the potential was www.annualreviews.org • A Life in Electrochemistry


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a

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Time (s) Figure 5 Principle of single nanoparticle collision experiments. (a) Scheme of single nanoparticle collision at the Au ultramicroelectrode UME surface. The reaction is switched on when the particle is in contact with the detection electrode and hydrazine oxidation is electrocatalyzed. (b) Current amplification: showing the difference between the hydrazine oxidation rate with Au and Pt UMEs. The arrow shows potential where the rate is fast on Pt and slow on Au. Scan rate, 50 mV/s; electrolyte, 10 mM of hydrazine + 50 mM of (phosphate buffered saline) PBS buffer; pH ∼7.5. (c) Representative current profile observed in a single nanoparticle collision event. From Reference 60, copyright 2008 by the American Chemical Society.

scanned back in a short time, its activity could be partially restored. Another case involved showing that a glass pipet pulled to a small (nm) orifice and immersed with the same electrolyte solution, e.g., KCl, on the inside and outside, would show rectification (i.e., a nonlinear current versus potential plot) when Ag/AgCl electrodes inside and outside the pipet were scanned, because the double layer at the pipet orifice affected the flow of ions (58). Later, this effect was found in what became known as nanopores. We didn’t think about giving the orifice a name, so the connection with these later experiments was largely lost. Our recent work has been directed to looking at electrode processes on a particle by particle (or eventually molecule by molecule) basis. This requires very low concentrations of electroactive particles that diffuse and migrate to an electrode such that one sees the result of a series of collisions, which show stochastic behavior (59). A representative figure and some references are given, but space does not permit a more extended treatment (Figure 5).

DISCLOSURE STATEMENT The author is not aware of any affiliations, memberships, funding, or financial holdings that might be perceived as affecting the objectivity of this review. 12.18

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ACKNOWLEDGMENTS The outstanding contributions of the many members of my group (graduate students, postdoctoral fellows, visitors) over the years have been responsible for essentially all of the work described. I did not try to list the names within the text, because I didn’t want to minimize the work of those that I didn’t discuss, but some of the names are shown in the references. (A complete list of publications is given at http://bard.cm.utexas.edu/styled-6/.) However I must acknowledge the work of two long-term members of my group, Fu-Ren Frank Fan and Chongyang Liu. Both are wonderful scientists who could do what at first seemed impossible experiments.


LITERATURE CITED 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12.

13. 14. 15. 16. 17. 18. 19. 20. 21. 22.

Stella A. 2013. Moissan’s furnace. Chem. World, June 3, p. 43 Jackson WM, ed. 1910. The Book of Knowledge. New York: Grolier, Inc. Innes WT. 1935. Exotic Aquarium Fishes. Philadelphia: Innes Publ. Co. Lingane JJ. 1953. Electroanalytical Chemistry. Olney, UK: Intersci. Publ. Kolthoff IM, Lingane JJ. 1952. Polarography. Hoboken, NJ: Wiley Delahay P. 1954. New Instrumental Methods in Electrochemistry. Olney, UK: Intersci. Publ. Geske DH, Bard AJ. 1959. Evaluation of the effect of secondary reactions in controlled potential coulometry. J. Phys. Chem. 63:1057–62 Bard AJ. 1961. Effect of electrode configuration and transition time in solid electrode chronopotentiometry. Anal. Chem. 33:11–15 Bard AJ. 1963. Correction for the inconstancy of the chronopotentiometric constant at short transition times. Anal. Chem. 35:340–43 Herman HB, Bard AJ. 1963. Cyclic chronopotentiometry. Diffusion controlled electrode reaction of a single component system. Anal. Chem. 35:1121–25 Bard AJ, Mayell JS. 1963. The electroreduction of quaternary ammonium compounds. J. Am. Chem. Soc. 85:421–25 Childs WV, Maloy JT, Keszthelyi CP, Bard AJ. 1971. Voltammetric and coulometric studies of the mechanism of electrohydrodimerization of diethyl fumarate in dimethylformadine solutions. J. Electrochem. Soc. 118:874–80 Bard AJ, Santhanam KSV, Maloy JT, Phelps J, Wheeler LO. 1968. Steric effects and the electrochemistry of phenyl-substituted anthracenes and related compounds. Disc. Faraday Soc. 45:167–74 Bard AJ. 1971. The electrochemistry of organic compounds in aprotic solvents—methods and applications. Pure Appl. Chem. 25:379–93 Santhanam KSV, Bard AJ. 1965. Chemiluminescence of electrogenerated 9,10-diphenylanthracene anion radical. J. Am. Chem. Soc. 87:139–140 Maloy JT, Prater KB, Bard AJ. 1968. Electrogenerated chemiluminescence. II. The rotating ring-disk electrode and the pyrene-N,N,N ,N -tetramethyl-p-phenylenediamine system. J. Phys. Chem. 72:4348–50 Faulkner LR, Bard AJ. 1969. Electrogenerated chemiluminescence. IV. Magnetic field effects on the electrogenerated chemiluminescence of some anthracenes. J. Am. Chem. Soc. 91:209–10 Cruser SA, Bard AJ. 1967. Concentration-intensity relationships in electrogenerated chemiluminescence. Anal. Lett. 1:11–17 Tokel NE, Bard AJ. 1972. Electrogenerated chemiluminescence. IX. Electrochemistry and emission from systems containing tris(2,2 -bipyridine)ruthenium(II) dichloride. J. Am. Chem. Soc. 94(8):2862–63 Rubinstein I, Bard AJ. 1981. Electrogenerated chemiluminescence. 37. Aqueous Ecl systems based on Ru(2,2 -bipyridine)3 2+ and oxalate or organic acids. J. Am. Chem. Soc. 103:512–16 Ege D, Becker WG, Bard AJ. 1984. Electrogenerated chemiluminescent determination of Ru(bpy)3 2+ at low levels. Anal. Chem. 56:2413–17 Hardee KL, Bard AJ. 1976. Semiconductor electrodes. V. The application of chemically vapor deposited iron oxide films to photosensitized electrolysis. J. Electrochem. Soc. 123:1024–26 www.annualreviews.org • A Life in Electrochemistry


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23. Hardee KL, Bard AJ. 1975. Semiconductor electrodes. I. The chemical vapor deposition and application of polycrystalline N-type titanium dioxide electrodes to the photosensitized electrolysis of water. J. Electrochem. Soc. 122:739–42 24. Frank SN, Bard AJ. 1977. Heterogeneous photocatalytic oxidation of cyanide ion in aqueous solutions at titanium dioxide powder. J. Am. Chem. Soc. 99:303–4 25. Kraeutler B, Bard AJ. 1977. Photoelectrosynthesis of ethane from acetate ion at an n-type TiO2 electrode. The photo-Kolbe reaction. J. Am. Chem. Soc. 99:7729–31 26. Kraeutler B, Bard AJ. 1978. Heterogeneous photocatalytic synthesis of methane from acetic acid—new Kolbe reaction pathway. J. Am. Chem. Soc. 100:2239–40 27. Dunn WW, Aikawa Y, Bard AJ. 1981. Heterogeneous photosynthetic production of amino acids at Pt/TiO2 suspensions by near ultraviolet light. J. Am. Chem. Soc. 103:6893–97 28. Frank SN, Bard AJ. 1975. Semiconductor electrodes. II. Electrochemistry at n-type TiO2 electrodes in acetonitrile solutions. J. Am. Chem. Soc. 97:7427–33 29. Bard AJ, Wrighton MS. 1977. Thermodynamic potential for the anodic dissolution of n-type semiconductors. A crucial factor controlling durability and efficiency in photochemical cells and an important criterion in the selection of new electrode/electrolyte systems. J. Electrochem. Soc. 124:1706–10 30. Krishnan M, White JR, Fox MA, Bard AJ. 1983. Integrated chemical systems: photocatalysis at semiconductors incorporated into polymer (Nafion)/mediator systems. J. Am. Chem. Soc. 105:7002–3 31. Saji T, Pasch NF, Webber SE, Bard AJ. 1978. Electrochemical behavior of polymers in aprotic media. 1. Polyvinylnaphthalene and polyvinylanthracene. J. Phys. Chem. 82:1101–5 32. Flanagan JB, Margel S, Bard AJ, Anson FC. 1978. Electron transfer to and from molecules containing multiple, noninteracting Redox centers. Electrochemical oxidation of poly(vinylferrocene). J. Am. Chem. Soc. 100:4248–53 33. Rubinstein I, Bard AJ. 1980. Polymer films on electrodes. IV. Nafion-coated electrodes and electrogenerated chemiluminescence of surface attached Ru(bpy)3 2+ . J. Am. Chem. Soc. 102:6641–42 34. Abruna HD, Bard AJ. 1981. Semiconductor electrodes. 40. Photoassisted hydrogen evolution at poly(benzyl viologen)-coated p-type silicon electrodes. J. Am. Chem. Soc. 103:6898–901 35. Henning TP, White HS, Bard AJ. 1982. Polymer films on electrodes. 10. Electrochemical behavior of solution species at Nafion-tetrathiafulvalenium bromide polymers. J. Am. Chem. Soc. 104:5862–68 36. Liu H-Y, Fan F-RF, Lin CW, Bard AJ. 1986. Scanning electrochemical and tunneling ultramicroelectrode microscope for high-resolution examination of electrode surfaces in solution. J. Am. Chem. Soc. 108:3838– 39 37. Bard AJ, Fan F-RF, Kwak J, Lev O. 1989. Scanning electrochemical microscopy. 1. Introduction and principles. Anal. Chem. 61:132–38 38. Bard AJ, Zoski CG. 2000. Voltammetry retrospective. Anal. Chem. 72:346A–52 39. Fan F-RF, Bard AJ. 1995. Electrochemical detection of single molecules. Science 267:871–74 40. Palacios RE, Fan F-RF, Bard AJ, Barbara PF. 2006. Single-molecule spectroelectrochemistry (SMS-EC). J. Am. Chem. Soc. 128:9028–29 41. Palacios RE, Fan F-RF, Grey JK, Suk J, Bard AJ, Barbara PF. 2007. Charging and discharging of single conjugated-polymer nanoparticles. Nat. Mater. 6:680–85 42. Fan F-RF, Kwak J, Bard AJ. 1996. Single molecule electrochemistry. J. Am. Chem. Soc. 118:9669–75 43. Lee J, Ye H, Pan S, Bard AJ. 2008. Screening of photocatalysts by scanning electrochemical microscopy. Anal. Chem. 80:7445–50 44. Chang H, Bard AJ. 1990. Formation of monolayer pits of controlled nanometer size on highly oriented pyrolytic graphite by gasification reactions as studied by scanning tunneling microscopy. J. Am. Chem. Soc. 112:4598–99 45. Hu K, Fan F-RF, Bard AJ, Hillier A. 1997. Direct measurement of diffuse double-layer forces at the semiconductor/electrolyte interface using an atomic force microscope. J. Phys. Chem. 101:8298–303 46. Gregg BA, Fox MA, Bard AJ. 1989. 2,3,7,8,12,13,17,18-octakis(β-hydroxyethyl)porphyrin (octaethanolporphyrin) and its liquid crystalline derivatives: synthesis and characterization. J. Am. Chem. Soc. 111:3024– 29 47. Gregg BA, Fox MA, Bard AJ. 1990. Photovoltaic effect in symmetrical cells of a liquid crystal porphyrin. J. Phys. Chem. 94:1586–98


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48. Liu C-Y, Pan H-L, Tang H, Fox MA, Bard AJ. 1995. Effect of structural order on the dark and photocurrents in zinc octakis (β-decoxyethyl)porphyrin thin-layer cells. J. Phys. Chem. 99:7632–36 49. Liu C-Y, Pan H-L, Fox MA, Bard AJ. 1997. Reversible charge trapping/detrapping in a photoconductive insulator of liquid crystal zinc porphyrin. Chem. Mater. 9:1422–29 50. Crooks RM, Fan F-RF, Bard AJ. 1984. Electrochemistry in near-critical and supercritical fluids. 1. Ammonia. J. Am. Chem. Soc. 106:6851–52 51. Tinker LA, Bard AJ. 1979. Electrochemistry in liquid sulfur dioxide. I. Oxidation of thianthrene, penothiazine, and 9,10-diphenylanthracene. J. Am. Chem. Soc. 101:2316–19 52. Flarsheim WM, Bard AJ, Johnston KP. 1989. Pronounced pressure effects on reversible electrode reactions in supercritical water. J. Phys. Chem. 93:4234–42 53. Teherani TH, Peer WJ, Lagowski JJ, Bard AJ. 1978. Anodic electrodeposition of gold from liquid ammonia solutions. J. Electrochem. Soc. 125:1717–18 54. Gaudiello JG, Sharp PR, Bard AJ. 1982. Electrochemistry in liquid sulfur dioxide. 3. The electrochemical production of new highly oxidized 2,2 -bipyridine complexes of ruthenium and iron. J. Am. Chem. Soc. 104:6367–77 55. Liu C-Y, Bard AJ. 2005. Electrochemistry and electrogenerated chemiluminescence with a single faradaic electrode. Anal. Chem. 77:5339–43 56. Liu C-Y, Bard AJ. 2008. Electrostatic electrochemistry at insulators. Nat. Mater. 7(6):505–9 57. Santhanam KSV, Jesperson N, Bard AJ. 1977. Application of a novel thermistor mercury electrode to the study of changes of activity of an adsorbed enzyme on electrochemical reduction and oxidation. J. Am. Chem. Soc. 99:274–76 58. Wei C, Bard AJ, Feldberg S. 1997. Current rectification at quartz nanopipet electrodes. Anal. Chem. 22:4627–33 59. Xiao XY, Bard AJ. 2007. Observing single nanoparticle collisions at an ultramicroelectrode by electrocatalytic amplification. J. Am. Chem. Soc. 129:9610–12 60. Xiao X, Fan F-RF, Zhou J, Bard AJ. 2008. Current transients in single nanoparticle collision events. J. Am. Chem. Soc. 130:16669–77



12.21

Electrochemistry of a Robust Neural Interface.

Electrosynthesis and electrochemistry.

Hydrogen-atom-mediated electrochemistry.

Electrochemistry in hollow-channel paper analytical devices.

Electrochemistry: Catalysis at the boundaries.

Single-Molecule Electrochemistry on a Porous Silica-Coated Electrode.

Oxygen electrochemistry as a cornerstone for sustainable energy conversion.

Electrochemistry of orthosilicate-based lithium battery cathodes: a perspective.

Electrochemistry and electrochemiluminescence from a redox-active metal-organic framework.

Electrochemistry and spectroelectrochemistry of bioactive hydroxyquinolines: a mechanistic study.

Ionic liquid electrolytes for reversible magnesium electrochemistry.

Nanoscale methods for single-molecule electrochemistry.

Eigenstress model for electrochemistry of solid surfaces.

Opportunities for microbial electrochemistry in municipal wastewater treatment--an overview.

Mesomorphism and electrochemistry of thienoviologen liquid crystals.

Electrochemistry of graphene and related materials.

mass spectrometry as a tool in metabolism studies-a review.

Solvents' Critical Role in Nonaqueous Lithium-Oxygen Battery Electrochemistry.

Electrochemistry in organisms. Electron flow and power output.

DNA Electrochemistry Shows DNMT1 Methyltransferase Hyperactivity in Colorectal Tumors.

Pyranopterin Coordination Controls Molybdenum Electrochemistry in Escherichia coli Nitrate Reductase.

Review of advances in coupling electrochemistry and liquid state NMR.

m-nitrobenzyl alcohol electrochemistry in fast atom bombardment mass spectrometry.

Electrochemistry of magnesium electrolytes in ionic liquids for secondary batteries.