Making raw materials for the manufacture of consumer goods produces high levels of carbon dioxide (CO2) emissions, involves hazardous materials, and requires high temperatures and pressures, usually generated by burning fossil fuels. MIT chemical engineers have now demonstrated a new approach that can operate on water plus electricity from renewable sources. Energized by a well-known catalyst, the process forms no CO2 emissions, requires no hazardous materials or extreme operating conditions, and generates just one byproduct—hydrogen. While much work remains, this new approach—relying on electricity and electrocatalysts—could one day significantly reduce the vast amounts of CO2 produced by the chemical industry today.
Most efforts to reduce energy consumption and carbon emissions have focused on the transportation and residential sectors. Little attention has been paid to industrial manufacturing, even though it consumes more energy than either of those sectors and emits high levels of CO2 in the process.
To help address that situation, Assistant Professor Karthish Manthiram, postdoc Kyoungsuk Jin, graduate students Joseph H. Maalouf and Minju Chung, and their colleagues, all of chemical engineering, have been devising new methods of synthesizing epoxides, a group of chemicals used in the manufacture of consumer goods ranging from polyester clothing, detergents, and antifreeze to pharmaceuticals and plastics.
“We don’t think about the embedded energy and carbon dioxide footprint of a plastic bottle we’re using or the clothing we’re putting on,” says Manthiram. “But epoxides are everywhere!”
As solar and wind and storage technologies mature, it’s time to address what Manthiram calls the “hidden energy and carbon footprints of materials made from epoxides.” And the key, he argues, may be to perform epoxide synthesis using electricity from renewable sources along with specially designed catalysts and an unlikely starting material: water.
Epoxides can be made from a variety of carbon-containing compounds known generically as olefins. But regardless of the olefin used, the conversion process generally produces high levels of CO2 or has other serious drawbacks.
To illustrate the problem, Manthiram describes processes now used to manufacture ethylene oxide, an epoxide used in making detergents, thickeners, solvents, plastics, and other consumer goods. Demand for ethylene oxide is so high that it has the fifth-largest CO2 footprint of any chemical made today.
The top panel in the figure below shows one common synthesis process. The recipe is simple: Combine ethylene molecules and oxygen molecules, subject the mixture to high temperatures and pressures, and separate out the ethylene oxide that forms.
However, as the diagram shows, those ethylene oxide molecules are accompanied by molecules of CO2—a problem, given the volume of ethylene oxide produced nationwide. In addition, the high temperatures and pressures required are generally produced by burning fossil fuels. And the conditions are so extreme that the reaction must take place in a massive pressure vessel. The capital investment required is high, so epoxides are generally produced in a central location and then transported long distances to the point of consumption.
Another widely synthesized epoxide is propylene oxide, which is used in making a variety of products, including perfumes, plasticizers, detergents, and polyurethanes. In this case, the olefin—propylene—is combined with tert-butyl hydroperoxide, as illustrated in the bottom panel above. An oxygen atom moves from the tert-butyl hydroperoxide molecule to the propylene to form the desired propylene oxide. The reaction conditions are somewhat less harsh than in ethylene oxide synthesis, but a side product must be dealt with. And while no CO2 is created, the tert-butyl hydroperoxide is highly reactive, flammable, and toxic, so it must be handled with extreme care.
In short, current methods of epoxide synthesis produce CO2, involve dangerous chemicals, require huge pressure vessels, or call for fossil fuel combustion. Manthiram and his team believed there must be a better way.
The goal in epoxide synthesis is straightforward: Simply transfer an oxygen atom from a source molecule onto an olefin molecule. Manthiram and his lab came up with an idea: Could water be used as a sustainable and benign source of the needed oxygen atoms? The concept was counterintuitive. “Organic chemists would say that it shouldn’t be possible because water and olefins don’t react with one another,” he says. “But what if we use electricity to liberate the oxygen atoms in water? Electrochemistry causes interesting things to happen—and it’s at the heart of what our group does.”
Using electricity to split water into oxygen and hydrogen is a standard practice called electrolysis. Usually, the goal of water electrolysis is to produce hydrogen gas for certain industrial applications or for use as a fuel. The oxygen is simply vented to the atmosphere.
To Manthiram, that practice seemed wasteful. Why not do something useful with the oxygen? Making an epoxide seemed the perfect opportunity—and the benefits could be significant. Generating two valuable products instead of one would bring down the high cost of water electrolysis. Indeed, it might become a cheaper, carbon-free alternative to today’s usual practice of producing hydrogen from natural gas. The electricity needed for the process could be generated from renewable sources such as solar and wind. There wouldn’t be any hazardous reactants or undesirable byproducts involved. And there would be no need for massive, costly, and accident-prone pressure vessels. As a result, epoxides could be made at small-scale, modular facilities close to the place they’re going to be used—no need to transport, distribute, or store the chemicals produced.
However, there was a chance that the proposed process might not work. During electrolysis, the oxygen atoms quickly pair up to form oxygen gas. The proposed process—illustrated in the diagram below—would require that some of the oxygen atoms move onto the olefin before they combine with one another.
To investigate the feasibility of the process, Manthiram’s group performed a fundamental analysis to find out whether the reaction is thermodynamically favorable. Does the energy of the overall system shift to a lower state by making the move? In other words, is the product more stable than the reactants were?
They started with a thermodynamic analysis of the proposed reaction at various combinations of temperature and pressure—the standard variables used in hydrocarbon processing. As an example, they again used ethylene oxide. The results, shown below, were not encouraging. As the uniform blue in the left-hand figure shows, even at elevated temperatures and pressures, the conversion of ethylene and water to ethylene oxide plus hydrogen doesn’t happen—just as a chemist’s intuition would predict.
But their proposal was to use voltage rather than pressure to drive the chemical reaction. As the right-hand figure above shows, with that change, the outcome of the analysis looked more promising. Conversion of ethylene to ethylene oxide occurs at around 0.8 volts. So the process is viable at voltages below that of an everyday AA battery and at essentially room temperature.
While a thermodynamic analysis can show that a reaction is possible, it doesn’t reveal how quickly it will occur, and reactions must be fast to be cost-effective. So the researchers needed to design a catalyst—a material that would speed up the reaction without getting consumed.
Designing catalysts for specific electrochemical reactions is a focus of Manthiram’s group. For this reaction, they decided to start with manganese oxide, a material known to catalyze the water-splitting reaction. And to increase the catalyst’s effectiveness, they fabricated it into nanoparticles—a particle size that would maximize the surface area on which reactions can take place.
The diagram below shows the special electrochemical cell they designed. Like all such cells, it has two electrodes—in this case, an anode where oxygen is transferred to make an olefin into an epoxide, and a cathode where hydrogen gas forms. The anode is made of carbon paper decorated with the nanoparticles of manganese oxide (shown in yellow). The cathode is made of platinum. Between the anode and the cathode is an electrolyte that ferries electrically charged ions between them. In this case, the electrolyte is a mixture of a solvent, water (the oxygen source), and the olefin.
The magnified views show what happens at the two electrodes. The right-hand view shows the olefin and water (H2O) molecules arriving at the anode surface. Encouraged by the catalyst, the water molecules break apart, sending two electrons (negatively charged particles, e–) into the anode and releasing two protons (positively charged hydrogen ions, H+) into the electrolyte. The leftover oxygen atom (O) joins the olefin molecule on the surface of the electrode, forming the desired epoxide molecule.
The two liberated electrons travel through the anode and around the external circuit (shown in red), where they pass through a power source—ideally, fueled by a renewable source such as wind or solar—and gain extra energy. When the two energized electrons reach the cathode, they join the two protons arriving in the electrolyte and—as shown in the left-hand magnified view—they form hydrogen gas (H2), which exits the top of the cell.
Experiments with that setup have been encouraging. Thus far, the work has involved an olefin called cyclooctene, a well-known molecule that’s been widely used by people studying oxidation reactions. “Ethylene and the like are structurally more important and need to be solved, but we’re developing a foundation on a well-known molecule just to get us started,” says Manthiram.
Results have already allayed a major concern. In one test, the researchers applied 3.8 volts across their mixture at room temperature, and after 4 hours, about half of the cyclooctene had converted into its epoxide counterpart, cyclooctene oxide. “So that result confirms that we can split water to make hydrogen and oxygen and then intercept the oxygen atoms so they move onto the olefin and convert it into an epoxide,” says Manthiram.
But how efficiently does the conversion happen? If this reaction is perfectly efficient, one oxygen atom will move onto an olefin for every two electrons that go into the anode. Thus, one epoxide molecule will form for each hydrogen molecule that forms. Using special equipment, the researchers counted the number of epoxide molecules formed for each pair of electrons passing through the external circuit to form hydrogen.
That analysis showed that their conversion efficiency was 30% of the maximum theoretical efficiency. “That’s because the electrons are also doing other reactions—maybe making oxygen, for instance, or oxidizing some of the solvent,” says Manthiram. “But for us, 30% is a remarkable number for a new reaction that was previously unknown. For that to be the first step, we’re very happy about it.”
Manthiram recognizes that the efficiency might need to be twice as high or even higher for the process to be commercially viable. “Techno-economics will ultimately guide where that number needs to be,” he says. “But I would say that the heart of our discoveries so far is the realization that there is a catalyst that can make this happen. That’s what has opened up everything that we’ve explored since the initial discovery.”
Manthiram is cautious not to overstate the potential implications of the work. “We know what the outcome is,” he says. “We put olefin in, and we get epoxide out.” But to optimize the conversion efficiency they need to know at a molecular level all the steps involved in that conversion. For example, does the electron transfer first by itself, or does it move with a proton at the same time? How does the catalyst bind the oxygen atom? And how does the oxygen atom transfer to the olefin on the surface of the catalyst?
According to Manthiram, he and his group have hypothesized a reaction sequence, and several analytical techniques have provided a “handful of observables” that support it. But he admits that there is much more theoretical and experimental work to do to develop and validate a detailed mechanism that they can use to guide the optimization process. And then there are practical considerations, such as how to extract the epoxides from the electrochemical cell and how to scale up production.
Manthiram believes that this work on epoxides is just “the tip of the iceberg” for his group. There are many other chemicals they might be able to make using voltage and specially designed catalysts. And while some attempts may not work, with each one they’ll learn more about how voltages and electrons and surfaces influence the outcome.
He and his team predict that the face of the chemical industry will change dramatically in the years to come. The need to reduce CO2 emissions and energy use is already pushing research on chemical manufacturing toward using electricity from renewable sources. And that electricity will increasingly be made at distributed sites. “If we have solar panels and wind turbines everywhere, why not do chemical synthesis close to where the power is generated, and make commercial products close to the communities that need them?” says Manthiram. The result will be a distributed, electrified, and decarbonized chemical industry—and a dramatic reduction in both energy use and CO2 emissions.
This research was supported by MIT’s Department of Chemical Engineering and by National Science Foundation Graduate Research Fellowships. Further information can be found in:
K. Jin, J.H. Maalouf, N. Lazouski, N. Corbin, D. Yang, and K. Manthiram. “Epoxidation of cyclooctene using water as the oxygen atom source at manganese oxide electrocatalysts.” Journal of the American Chemical Society, vol. 141, pp. 6413–6418, 2019. Online: doi.org/10.1021/jacs.9b02345.
Z.J. Schiffer and K. Manthiram. “Electrification and decarbonization of the chemical industry.” Joule, vol. 1, pp. 10–14, September 6, 2017. Online: doi.org/10.1016/j.joule.2017.07.008.