Making a light-harvesting antenna from scratch: Biomimetic antenna for gathering sunlight may one day transform solar-powered devices

Sometimes when people talk about solar energy, they tacitly assume that we're stuck with some version of the silicon solar cell and its technical and cost limitations. Not so.

The invention of the solar cell, in 1941, was inspired by a newfound understanding of semiconductors, materials that can use light energy to create mobile electrons -- and ultimately an electrical current.

Silicon solar cells have almost nothing to do with the biological photosystems in tree leaves and pond scum that use light energy to push electrons across a membrane -- and ultimately create sugars and other organic molecules.

At the time, nobody understood these complex assemblages of proteins and pigments well enough to exploit their secrets for the design of solar cells.

But things have changed.

At Washington University in St. Louis's Photosynthetic Antenna Research Center (PARC) scientists are exploring native biological photosystems, building hybrids that combine natural and synthetic parts, and building fully synthetic analogs of natural systems.

One team has just succeeded in making a crucial photosystem component -- a light-harvesting antenna -- from scratch. The new antenna is modeled on the chlorosome found in green bacteria.

Chlorosomes are giant assemblies of pigment molecules. Perhaps Nature's most spectacular light-harvesting antennae, they allow green bacteria to photosynthesize even in the dim light in ocean deeps.

Dewey Holten, PhD, professor of chemistry in Arts & Sciences, ard collaborator Christine Kirmaier, PhD, research professor of chemistry are part of a team that is trying to make synthetic chlorosomes. Holten and Kirmaier use ultra-fast laser spectroscopy and other analytic techniques to follow the rapid-fire energy transfers in photosynthesis.

His team's latest results were described in a recent issue of New Journal of Chemistry.

Chlorosomes

Biological systems that capture the energy in sunlight and convert it to the energy of chemical bonds come in many varieties, but they all have two basic parts: the light harvesting complexes, or antennae, and the reaction center complexes. The antennae consist of many pigment molecules that absorb photons and pass the excitation energy to the reaction centers.

In the reaction centers, the excitation energy sets off a chain of reactions that create ATP, a molecule often called the energy currency of the cell because the energy stored ATP powers most cellular work. Cellular organelles selectively break those bonds in ATP molecules when they need an energy hit for cellular work.

Green bacteria, which live in the lower layers of ponds, lakes and marine environments, and in the surface layers of sediments, have evolved large and efficient light-harvesting antennae very different from those found in plants bathing in sunlight on Earth's surface.

The antennae consist of highly organized three-dimensional systems of as many as 250,000 pigment molecules that absorb light and funnel the light energy through a pigment/protein complex called a baseplate to a reaction center, where it triggers chemical reactions that ultimately produce ATP.

In plants and algae (and in the baseplate in the green bacteria) photo pigments are bound to protein scaffolds, which space and orient the pigment molecules in such a way that energy is efficiently transferred between them.

But chlorosomes don't have a protein scaffold. Instead the pigment molecules self -assemble into a structure that supports the rapid migration of excitation energy.

This is intriguing because it suggests chlorosome mimics might be easier to incorporate in the design of solar devices than biomimetics that are made of proteins as well as pigments.

Synthetic pigments

The goal of the work described in the latest journal article was to see whether synthesized pigment molecules could be induced to self-assemble. The process by which the pigments align and bond is not well understood.

"The structure of the pigment assemblies in chlorosomes is the subject of intense debate," Holten says, "and there are several competing models for it."

Given this uncertainty, the scientists wanted to study many variations of a pigment molecule to see what favored and what blocked assembly.

A chemist wishing to design pigments that mimic those found in photosynthetic organisms first builds one of three molecular frameworks. All three are macrocycles, or giant rings: porphyrin, chlorin and bacteriochlorin.

"One of the members of our team, Jon Lindsey can synthesize analogs of all three pigment types from scratch," says Holten. (Lindsey, PhD, is Glaxo Professor of Chemistry at North Carolina State University.)

In the past, chemists making photo pigments have usually started with porphyrins, which are the easiest of the three types of macrocycles to synthesize. But Lindsey also has developed the means to synthesize chlorins, the basis for the pigments found in the chlorosomes of green bacteria. The chlorins push the absorption to the red end of the visible spectrum, an area of the spectrum scientists would like to be able to harvest for energy.

Key to pigment self-assembly are the metal atoms and hydroxyl (OH) and carbonyl (C=O) groups in the pigment molecules (the groups shown in color in the above illustration).

Doctoral student Olga Mass and coworkers in Lindsey's lab synthesized 30 different chlorins, systematically adding or removing chemical groups thought to be important for self-assembly but also attaching peripheral chemical groups that take up space and might make it harder for the molecules to stack or that shift around the distributions of electrons so that the molecules might stack more easily.

Testing for aggregation

The powdered pigments were carefully packaged and shipped by Fed Ex (because the Post Office won't ship chemicals) to Holten's lab at WUSTL and to David Bocian's lab at the University of California at Riverside.

Scientists in both labs made up green-tinctured solutions of each of the 30 molecules in small test tubes and then poked and prodded the solutions by means of analytical techniques to see whether the pigment had aggregated and, if so, how much had formed the assemblies. Holten's lab studied their absorption of light and their fluorescence (which indicated the presence of monomers, since assemblies don't normally fluoresce) and Bocian's lab studied their vibrational properties, which are determined by the network of bonds in the molecule or pigment aggregate as a whole.

In one crucial test Joseph Springer, a PhD student in Holten's lab, compared the absorption spectrum of a pigment in a polar solvent that would prevent it from self-assembling to the spectrum of the pigment in a nonpolar solvent that would allow the molecules to interact with one another and form assemblies.

"You can see them aggregate," Springer says. "A pigment that is totally in solution is clear, but colored a brilliant green. When it aggregates, the solution becomes a duller green and you can see tiny flecks in the liquid."

The absorption spectra indicated that some pigments formed extensive assemblies and that the steric and electronic properties of the molecules predicted the degree to which they would assemble.

Up next

Although this project focused on self-assembly, the PARC scientists have already taken the next step toward a practical solar device. "With Pratim Biswas, PhD, the Lucy and Stanley Lopata Professor and chair of the Department of Energy, Environmental & Chemical Engineering, we've since demonstrated that we can get the pigments to self-assemble on surfaces, which is the next step in using them to design solar devices," says Holten.

"We're not trying to make a more efficient solar cell in the next six months," Holten cautions. "Our goal instead is to develop fundamental understanding so that we can enable the next generation of more efficient solar powered devices."

Biomimicry hasn't always worked. Engineers often point out early flying machines that attempted to mimic birds didn't work and that flying machines stayed aloft only when nventors abandoned biological models and came up with their own designs.

But there is nothing predestined or inevitable about this. As biological knowledge has exploded in the past 50 years, mimicking nature has become a smarter strategy. Biomimetic or biohybrid designs already have solved significant engineering problems in other areas and promise to greatly improve the design of solar powered devices as well.

After all, Nature has had billions of years to experiment with ways to harness the energy in sunlight for useful work.

Story Source:

The above story is reprinted from materials provided by Washington University in St. Louis. The original article was written by Diana Lutz.

Note: Materials may be edited for content and length. For further information, please contact the source cited above.

Journal Reference:

Olga Mass, Dinesh R. Pandithavidana, Marcin Ptaszek, Koraliz Santiago, Joseph W. Springer, Jieying Jiao, Qun Tang, Christine Kirmaier, David F. Bocian, Dewey Holten, Jonathan S. Lindsey. De novo synthesis and properties of analogues of the self-assembling chlorosomal bacteriochlorophylls. New Journal of Chemistry, 2011; 35 (11): 2671 DOI: 10.1039/C1NJ20611G

'Artificial leaf' makes fuel from sunlight (w/ video)

Researchers led by MIT professor Daniel Nocera have produced something they’re calling an “artificial leaf”: Like living leaves, the device can turn the energy of sunlight directly into a chemical fuel that can be stored and used later as an energy source.

The artificial leaf — a silicon solar cell with different catalytic materials bonded onto its two sides — needs no external wires or control circuits to operate. Simply placed in a container of water and exposed to sunlight, it quickly begins to generate streams of bubbles: oxygen bubbles from one side and hydrogen bubbles from the other. If placed in a container that has a barrier to separate the two sides, the two streams of bubbles can be collected and stored, and used later to deliver power: for example, by feeding them into a fuel cell that combines them once again into water while delivering an electric current.

The creation of the device is described in a paper published Sept. 30 in the journal Science. Nocera, the Henry Dreyfus Professor of Energy and professor of chemistry at MIT, is the senior author; the paper was co-authored by his former student Steven Reece PhD ’07 (who now works at Sun Catalytix, a company started by Nocera to commercialize his solar-energy inventions), along with five other researchers from Sun Catalytix and MIT.

The device, Nocera explains, is made entirely of earth-abundant, inexpensive materials — mostly silicon, cobalt and nickel — and works in ordinary water. Other attempts to produce devices that could use sunlight to split water have relied on corrosive solutions or on relatively rare and expensive materials such as platinum.

The artificial leaf is a thin sheet of semiconducting silicon — the material most solar cells are made of — which turns the energy of sunlight into a flow of wireless electricity within the sheet. Bound onto the silicon is a layer of a cobalt-based catalyst, which releases oxygen, a material whose potential for generating fuel from sunlight was discovered by Nocera and his co-authors in 2008. The other side of the silicon sheet is coated with a layer of a nickel-molybdenum-zinc alloy, which releases hydrogen from the water molecules.

This video is not supported by your browser at this time.

An 'artificial leaf' made by Daniel Nocera and his team, using a silicon solar cell with novel catalyst materials bonded to its two sides, is shown in a container of water with light (simulating sunlight) shining on it. The light generates a flow of electricity that causes the water molecules, with the help of the catalysts, to split into oxygen and hydrogen, which bubble up from the two surfaces.Video courtesy of the Nocera Lab/Sun Catalytix

“I think there’s going to be real opportunities for this idea,” Nocera says. “You can’t get more portable — you don’t need wires, it’s lightweight,” and it doesn’t require much in the way of additional equipment, other than a way of catching and storing the gases that bubble off. “You just drop it in a glass of water, and it starts splitting it,” he says.

Now that the “leaf” has been demonstrated, Nocera suggests one possible further development: tiny particles made of these materials that can split water molecules when placed in sunlight — making them more like photosynthetic algae than leaves. The advantage of that, he says, is that the small particles would have much more surface area exposed to sunlight and the water, allowing them to harness the sun’s energy more efficiently. (On the other hand, engineering a system to separate and collect the two gases would be more complicated in such a setup.)

The new device is not yet ready for commercial production, since systems to collect, store and use the gases remain to be developed. “It’s a step,” Nocera says. “It’s heading in the right direction.”

Ultimately, he sees a future in which individual homes could be equipped with solar-collection systems based on this principle: Panels on the roof could use sunlight to produce hydrogen and oxygen that would be stored in tanks, and then fed to a fuel cell whenever electricity is needed. Such systems, Nocera hopes, could be made simple and inexpensive enough so that they could be widely adopted throughout the world, including many areas that do not presently have access to reliable sources of electricity.

Professor James Barber, a biochemist from Imperial College London who was not involved in this research, says Nocera’s 2008 finding of the cobalt-based catalyst was a “major discovery,” and these latest findings “are equally as important, since now the water-splitting reaction is powered entirely by visible light using tightly coupled systems comparable with that used in natural photosynthesis. This is a major achievement, which is one more step toward developing cheap and robust technology to harvest solar energy as chemical fuel.”

Barber cautions that “there will be much work required to optimize the system, particularly in relation to the basic problem of efficiently using protons generated from the water-splitting reaction for hydrogen production.” But, he says, “there is no doubt that their achievement is a major breakthrough which will have a significant impact on the work of others dedicated to constructing light-driven catalytic systems to produce hydrogen and other solar fuels from water. This technology will advance side by side with new initiatives to improve and lower the cost of photovoltaics.”

Nocera’s ongoing research with the artificial leaf is directed toward “driving costs lower and lower,” he says, and looking at ways of improving the system’s efficiency. At present, the leaf can redirect about 2.5 percent of the energy of sunlight into hydrogen production in its wireless form; a variation using wires to connect the catalysts to the solar cell rather than bonding them together has attained 4.7 percent efficiency. (Typical commercial solar cells today have efficiencies of more than 10 percent). One question Nocera and his colleagues will be addressing is which of these configurations will be more efficient and cost-effective in the long run.

Another line of research is to explore the use of photovoltaic (solar cell) materials other than silicon — such as iron oxide, which might be even cheaper to produce. “It’s all about providing options for how you go about this,” Nocera says.

More information: http://www.science … 816.full.pdf

Provided by Massachusetts Institute of Technology (news : web)

Novel alloy could produce hydrogen fuel from sunlight

Scientists from the University of Kentucky and the University of Louisville have determined that an inexpensive semiconductor material can be "tweaked" to generate hydrogen from water using sunlight.

The research, funded by the U.S. Department of Energy, was led by Professors Madhu Menon and R. Michael Sheetz at the UK Center for Computational Sciences, and Professor Mahendra Sunkara and graduate student Chandrashekhar Pendyala at the UofL Conn Center for Renewable Energy Research. Their findings were published Aug. 1 in the Physical Review Journal (Phys Rev B 84, 075304).

The researchers say their findings are a triumph for computational sciences, one that could potentially have profound implications for the future of solar energy.

Using state-of-the-art theoretical computations, the UK-UofL team demonstrated that an alloy formed by a 2 percent substitution of antimony (Sb) in gallium nitride (GaN) has the right electrical properties to enable solar light energy to split water molecules into hydrogen and oxygen, a process known as photoelectrochemical (PEC) water splitting. When the alloy is immersed in water and exposed to sunlight, the chemical bond between the hydrogen and oxygen molecules in water is broken. The hydrogen can then be collected.

"Previous research on PEC has focused on complex materials," Menon said. "We decided to go against the conventional wisdom and start with some easy-to-produce materials, even if they lacked the right arrangement of electrons to meet PEC criteria. Our goal was to see if a minimal 'tweaking' of the electronic arrangement in these materials would accomplish the desired results."

Gallium nitride is a semiconductor that has been in widespread use to make bright-light LEDs since the 1990s. Antimony is a metalloid element that has been in increased demand in recent years for applications in microelectronics. The GaN-Sb alloy is the first simple, easy-to-produce material to be considered a candidate for PEC water splitting. The alloy functions as a catalyst in the PEC reaction, meaning that it is not consumed and may be reused indefinitely. UofL and UK researchers are currently working toward producing the alloy and testing its ability to convert solar energy to hydrogen.

Hydrogen has long been touted as a likely key component in the transition to cleaner energy sources. It can be used in fuel cells to generate electricity, burned to produce heat, and utilized in internal-combustion engines to power vehicles. When combusted, hydrogen combines with oxygen to form water vapor as its only waste product. Hydrogen also has wide-ranging applications in science and industry.

Because pure hydrogen gas is not found in free abundance on Earth, it must be manufactured by unlocking it from other compounds. Thus, hydrogen is not considered an energy source, but rather an "energy carrier." Currently, it takes a large amount of electricity to generate hydrogen by water splitting. As a consequence, most of the hydrogen manufactured today is derived from non-renewable sources such as coal and natural gas.

Sunkara says the GaN-Sb alloy has the potential to convert solar energy into an economical, carbon-free source for hydrogen.

"Hydrogen production now involves a large amount of CO2 emissions," Sunkara said. "Once this alloy material is widely available, it could conceivably be used to make zero-emissions fuel for powering homes and cars and to heat homes."

Menon says the research should attract the interest of other scientists across a variety of disciplines.

"Photocatalysis is currently one of the hottest topics in science," Menon said. "We expect the present work to have a wide appeal in the community spanning chemistry, physics and engineering."

Provided by University of Kentucky

Novel alloy could produce hydrogen fuel from sunlight

Scientists from the University of Kentucky and the University of Louisville have determined that an inexpensive semiconductor material can be "tweaked" to generate hydrogen from water using sunlight.

The research, funded by the U.S. Department of Energy, was led by Professors Madhu Menon and R. Michael Sheetz at the University of Kentucky Center for Computational Sciences, and Professor Mahendra Sunkara and graduate student Chandrashekhar Pendyala at the University of Louisville Conn Center for Renewable Energy Research. Their findings were published Aug. 1 in the Physical Review B.

The researchers say their findings are a triumph for computational sciences, one that could potentially have profound implications for the future of solar energy.

Using state-of-the-art theoretical computations, the University of Kentucky-University of Louisville team demonstrated that an alloy formed by a 2 percent substitution of antimony (Sb) in gallium nitride (GaN) has the right electrical properties to enable solar light energy to split water molecules into hydrogen and oxygen, a process known as photoelectrochemical (PEC) water splitting. When the alloy is immersed in water and exposed to sunlight, the chemical bond between the hydrogen and oxygen molecules in water is broken. The hydrogen can then be collected.

"Previous research on PEC has focused on complex materials," Menon said. "We decided to go against the conventional wisdom and start with some easy-to-produce materials, even if they lacked the right arrangement of electrons to meet PEC criteria. Our goal was to see if a minimal 'tweaking' of the electronic arrangement in these materials would accomplish the desired results."

Gallium nitride is a semiconductor that has been in widespread use to make bright-light LEDs since the 1990s. Antimony is a metalloid element that has been in increased demand in recent years for applications in microelectronics. The GaN-Sb alloy is the first simple, easy-to-produce material to be considered a candidate for PEC water splitting. The alloy functions as a catalyst in the PEC reaction, meaning that it is not consumed and may be reused indefinitely. University of Louisville and University of Kentucky researchers are currently working toward producing the alloy and testing its ability to convert solar energy to hydrogen.

Hydrogen has long been touted as a likely key component in the transition to cleaner energy sources. It can be used in fuel cells to generate electricity, burned to produce heat, and utilized in internal-combustion engines to power vehicles. When combusted, hydrogen combines with oxygen to form water vapor as its only waste product. Hydrogen also has wide-ranging applications in science and industry.

Because pure hydrogen gas is not found in free abundance on Earth, it must be manufactured by unlocking it from other compounds. Thus, hydrogen is not considered an energy source, but rather an "energy carrier." Currently, it takes a large amount of electricity to generate hydrogen by water splitting. As a consequence, most of the hydrogen manufactured today is derived from non-renewable sources such as coal and natural gas.

Sunkara says the GaN-Sb alloy has the potential to convert solar energy into an economical, carbon-free source for hydrogen.

"Hydrogen production now involves a large amount of CO2 emissions," Sunkara said. "Once this alloy material is widely available, it could conceivably be used to make zero-emissions fuel for powering homes and cars and to heat homes."

Menon says the research should attract the interest of other scientists across a variety of disciplines.

"Photocatalysis is currently one of the hottest topics in science," Menon said. "We expect the present work to have a wide appeal in the community spanning chemistry, physics and engineering."

Story Source:

The above story is reprinted (with editorial adaptations ) from materials provided by University of Kentucky, via EurekAlert!, a service of AAAS.

Journal Reference:

R. Sheetz, Ernst Richter, Antonis N. Andriotis, Sergey Lisenkov, Chandrashekhar Pendyala, Mahendra K. Sunkara, Madhu Menon. Visible-light absorption and large band-gap bowing of GaN_{1-x}Sb_{x} from first principles. Physical Review B, 2011; 84 (7) DOI: 10.1103/PhysRevB.84.075304

Chemistry with sunlight: Combining electrochemistry and photovoltaics to clean up oxidation reactions

The idea is simple, says Kevin Moeller, PhD, and yet it has huge implications. All we are recommending is using photovoltaic cells (clean energy) to power electrochemical reactions (clean chemistry). Moeller is the first to admit this isn't new science.

"But we hope to change the way people do this kind of chemistry by making a connection for them between two existing technologies," he says.

To underscore the simplicity of the idea, Moeller and his co-authors used a $6 solar cell sold on the Internet and intended to power toy cars to run reactions described in an article published in Green Chemistry.

If their suggestion were widely adopted by the chemical industry, it would eliminate the toxic byproducts currently produced by a class of reactions commonly used in chemical synthesis — and with them the environmental and economic damage they cause.

The trouble with oxidation reactions

Moeller, a professor of chemistry in Arts & Sciences at Washington University in St. Louis, is an organic chemist, who makes and manipulates molecules made mainly of carbon, hydrogen, oxygen and nitrogen.
One important tool for synthesizing organic molecules — an enormous category that includes everything from anesthetics to yarn — is the oxidation reaction.

"They are the one tool we have that allows us to increase the functionality of a molecule, to add more "handles" to it by which it can be manipulated," says Moeller.

"Molecules interact with each other through combinations of atoms known as functional groups," he explains. "Ketones, alcohols or amines are all functional groups. The more functional groups you have on a molecule, the more you can control how the molecule interacts with others."

"Oxidation reactions attach functional groups to a molecule," he continues. "If I have a hydrocarbon that consists of nothing but carbon and hydrogen atoms bonded together, and I want to convert it to an alcohol, a ketone or an amine, I have to oxidize it."

In an oxidation reaction, an electron is removed from a molecule. But that electron has to go somewhere, so every oxidation reaction is paired with a reduction reaction, where an electron is added to a second molecule.

The problem, says Moeller, is that "that second molecule is a waste product; it's not something you want."

One example, he says, is an industrial alcohol oxidation that uses the oxidant chromium to convert an alcohol into a ketone. In the process the chromium, originally chromium VI, picks up electrons and becomes chromium IV. Chromium IV is the waste product of the oxidation reaction.

In this case, there is a partial solution. Sodium periodate is used to recycle the highly toxic chromium IV. A salt, the sodium periodate dissociates in solution and the periodate ion (an iodine atom with attached oxygens) interacts with the chromium, restoring it to its original oxidation state.

The catch is that restoring the chromium destroys the periodate. In addition, the process is inefficient; three equivalents of periodate is consumed for every equivalent of desired product produced.

Seeking cleaner byproducts

"All chemical oxidations have a byproduct, says Moeller, so the question is not whether there will be a byproduct but what that byproduct will be. People have starting thinking about how they might run oxidations where the reduced byproduct is something benign."

"If you use oxygen to do the oxidation, the byproduct is water, and that is a gentle process," he says.

But there is a catch. Like all other molecules, oxygen has a set oxidation potential, or willingness to accept electrons. "So whatever I want to oxidize in solution has to have an oxidation potential that matches oxygen's. If it doesn't, I might have to change my whole reaction around to make sure I can use oxygen. And when I change the whole reaction around, maybe it doesn't run as well as it used to. So I'm limited in what I can do," Moeller says.

Enlarge

?Electrochemistry can oxidize molecules with any oxidation potential, because the electrode voltage can be tuned or adjusted, or I can run the reaction in such a way that it adjusts itself. So I have tremendous versatility for doing things,? says Kevin Moeller, professor of chemistry at Washington University in St. Louis. Moreover, the byproduct of electrochemical oxidation is hydrogen gas, so this too is a clean process.Bit electrochemistry can be only as green as the source of the electricity. The answer is to use the cleanest possible energy, solar energy captured by photovoltaic cells, to run electrochemical reactions. Credit: David Kilper/WUSTL

There's another way to do it. "Electrochemistry can oxidize molecules with any oxidation potential, because the electrode voltage can be tuned or adjusted, or I can run the reaction in such a way that it adjusts itself. So I have tremendous versatility for doing things," says Moeller.

Moreover, the byproduct of electrochemical oxidation is hydrogen gas, so this too is a clean process.

But again there is a catch. Electrochemistry can be only as green as the source of the electricity. If the oxidation reaction is running clean, but the electricity comes from a coal-fired plant, the problem has not been avoided, just displaced.

The answer is to use the cleanest possible energy, solar energy captured by photovoltaic cells, to run electrochemical reactions.

"That's what the Green Chemistry article is about," says Moeller. "It's a proof-of-principle paper that says it's easy to make this work, and it works just like reactions that don't use photovoltaics, so the chemical reaction doesn't have to be changed around."

The next step

The Green Chemistry article demonstrated the method by directly oxidizing molecules at the electrode. No chemical reagent was used. Since writing the article, Moeller's group has been studying how solar-powered electrochemistry might be used to recycle chemical oxidants in a clean way.

Why would manufacturers choose to use a chemical oxidant, if the voltage of the electrode can be matched to the oxidation potential of the molecule that must be oxidized?

"An electrode selects purely on oxidation potential," Moeller explains. "A chemical reagent does not. The binding properties of the chemical reagent might differ from one part of the molecule to another. And there's also something called steric hindrance, which means that one part of the molecule might physically block access to an oxidation site, forcing substrates to other sites on the reagent."

"The chemistry community has learned how to use chemical reagents to make reactions selective," he says. "The reagents are usually expensive and toxic, so they are recycled," he says. "We are working on cleaning up reagent recycling."

In the chromium oxidation described above, for example, chromium IV could be recycled electrochemically instead of through a reaction with periodate. Instead of periodate waste[consistent with description above where periodate consumed?], the reaction would produce hydrogen gas as the byproduct.

"Another example is an industrial process for carrying out alcohol oxidations that convert the alcohol group to a carbonyl group," says Moeller. This process uses TEMPO, a complex chemical reagent discovered in 1960. TEMPO is expensive so it is recycled by the addition of bleach. This regenerates the TEMPO but produces sodium chloride as a byproduct."

In small quantities sodium chloride is table salt, but in industrial quantities it is a waste product whose disposal is costly. Once again, the TEMPO can be recycled using electrochemistry, a process that produces hydrogen as the only byproduct.

"We can't make all of chemical synthesis cleaner by hitching solar power to electrochemistry," Moeller says, "but we can fix the oxidation reactions that people use. And maybe that will inspire someone else to come up with simple and innovative solutions to other types of reactions they're interested in."

Provided by Washington University in St. Louis (news : web)

Chemistry with sunlight: Combining electrochemistry and photovoltaics to clean up oxidation reactions

The idea is simple, says Kevin Moeller, PhD, and yet it has huge implications. All we are recommending is using photovoltaic cells (clean energy) to power electrochemical reactions (clean chemistry). Moeller is the first to admit this isn't new science.

"But we hope to change the way people do this kind of chemistry by making a connection for them between two existing technologies," he says.

To underscore the simplicity of the idea, Moeller and his co-authors used a $6 solar cell sold on the Internet and intended to power toy cars to run reactions described in an article published in Green Chemistry.

If their suggestion were widely adopted by the chemical industry, it would eliminate the toxic byproducts currently produced by a class of reactions commonly used in chemical synthesis -- and with them the environmental and economic damage they cause.

The trouble with oxidation reactions

Moeller, a professor of chemistry in Arts & Sciences at Washington University in St. Louis, is an organic chemist, who makes and manipulates molecules made mainly of carbon, hydrogen, oxygen and nitrogen.

One important tool for synthesizing organic molecules -- an enormous category that includes everything from anesthetics to yarn -- is the oxidation reaction.

"They are the one tool we have that allows us to increase the functionality of a molecule, to add more "handles" to it by which it can be manipulated," says Moeller.

"Molecules interact with each other through combinations of atoms known as functional groups," he explains. "Ketones, alcohols or amines are all functional groups. The more functional groups you have on a molecule, the more you can control how the molecule interacts with others."

"Oxidation reactions attach functional groups to a molecule," he continues. "If I have a hydrocarbon that consists of nothing but carbon and hydrogen atoms bonded together, and I want to convert it to an alcohol, a ketone or an amine, I have to oxidize it."

In an oxidation reaction, an electron is removed from a molecule. But that electron has to go somewhere, so every oxidation reaction is paired with a reduction reaction, where an electron is added to a second molecule.

The problem, says Moeller, is that "that second molecule is a waste product; it's not something you want."

One example, he says, is an industrial alcohol oxidation that uses the oxidant chromium to convert an alcohol into a ketone. In the process the chromium, originally chromium VI, picks up electrons and becomes chromium IV. Chromium IV is the waste product of the oxidation reaction.

In this case, there is a partial solution. Sodium periodate is used to recycle the highly toxic chromium IV. A salt, the sodium periodate dissociates in solution and the periodate ion (an iodine atom with attached oxygens) interacts with the chromium, restoring it to its original oxidation state.

The catch is that restoring the chromium destroys the periodate. In addition, the process is inefficient; three equivalents of periodate is consumed for every equivalent of desired product produced.

Seeking cleaner byproducts

"All chemical oxidations have a byproduct, says Moeller, so the question is not whether there will be a byproduct but what that byproduct will be. People have starting thinking about how they might run oxidations where the reduced byproduct is something benign."

"If you use oxygen to do the oxidation, the byproduct is water, and that is a gentle process," he says.

But there is a catch. Like all other molecules, oxygen has a set oxidation potential, or willingness to accept electrons. "So whatever I want to oxidize in solution has to have an oxidation potential that matches oxygen's. If it doesn't, I might have to change my whole reaction around to make sure I can use oxygen. And when I change the whole reaction around, maybe it doesn't run as well as it used to. So I'm limited in what I can do," Moeller says.

A simpler idea is also cleaner.

There's another way to do it. "Electrochemistry can oxidize molecules with any oxidation potential, because the electrode voltage can be tuned or adjusted, or I can run the reaction in such a way that it adjusts itself. So I have tremendous versatility for doing things," says Moeller.

Moreover, the byproduct of electrochemical oxidation is hydrogen gas, so this too is a clean process.

But again there is a catch. Electrochemistry can be only as green as the source of the electricity. If the oxidation reaction is running clean, but the electricity comes from a coal-fired plant, the problem has not been avoided, just displaced.

The answer is to use the cleanest possible energy, solar energy captured by photovoltaic cells, to run electrochemical reactions.

"That's what the Green Chemistry article is about," says Moeller. "It's a proof-of-principle paper that says it's easy to make this work, and it works just like reactions that don't use photovoltaics, so the chemical reaction doesn't have to be changed around."

The next step

The Green Chemistry article demonstrated the method by directly oxidizing molecules at the electrode. No chemical reagent was used. Since writing the article, Moeller's group has been studying how solar-powered electrochemistry might be used to recycle chemical oxidants in a clean way.

Why would manufacturers choose to use a chemical oxidant, if the voltage of the electrode can be matched to the oxidation potential of the molecule that must be oxidized?

"An electrode selects purely on oxidation potential," Moeller explains. "A chemical reagent does not. The binding properties of the chemical reagent might differ from one part of the molecule to another. And there's also something called steric hindrance, which means that one part of the molecule might physically block access to an oxidation site, forcing substrates to other sites on the reagent."

"The chemistry community has learned how to use chemical reagents to make reactions selective," he says. "The reagents are usually expensive and toxic, so they are recycled," he says. "We are working on cleaning up reagent recycling."

In the chromium oxidation described above, for example, chromium IV could be recycled electrochemically instead of through a reaction with periodate. Instead of periodate waste[consistent with description above where periodate consumed?], the reaction would produce hydrogen gas as the byproduct.

"Another example is an industrial process for carrying out alcohol oxidations that convert the alcohol group to a carbonyl group," says Moeller. This process uses TEMPO, a complex chemical reagent discovered in 1960. TEMPO is expensive so it is recycled by the addition of bleach. This regenerates the TEMPO but produces sodium chloride as a byproduct."

In small quantities sodium chloride is table salt, but in industrial quantities it is a waste product whose disposal is costly. Once again, the TEMPO can be recycled using electrochemistry, a process that produces hydrogen as the only byproduct.

"We can't make all of chemical synthesis cleaner by hitching solar power to electrochemistry," Moeller says, "but we can fix the oxidation reactions that people use. And maybe that will inspire someone else to come up with simple and innovative solutions to other types of reactions they're interested in."

Story Source:

The above story is reprinted (with editorial adaptations by ScienceDaily staff) from materials provided by Washington University in St. Louis.

Journal Reference:

Laura A. Anderson, Alison Redden, Kevin D. Moeller. Connecting the dots: using sunlight to drive electrochemical oxidations. Green Chemistry, 2011; DOI: 10.1039/C1GC15207F

Note: If no author is given, the source is cited instead.

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