 Good afternoon and welcome to today's Energy Seminar. Again, we have a real treat this quarter. You may or may not know that Stanford has this tremendously competitive competition for the best student and postdoc energy research talks. Goes on all summer, I think you have to be nominated to get in it. There's about 20 people in it. Only four people win and you're going to hear from two of them today. So they were worried where you are now and not that long ago, and maybe where you are right now. But to introduce them right, I'd like to introduce Jenny Milne, who is the Associate Director of the Precourt Institute for Advanced Research Projects, who actually ran the competition with a couple of colleagues. But also you heard about last week from the accelerator person. She was the one that they said one of two people who actually enabled the group to actually review 100 faculty proposals and make decisions on 30 selected projects in a lot less time than they thought they were going to have. Jenny is one of the main people that pulled that up. So Jenny, please take that floor. Thank you, John, very much for that very nice introduction. So thank you all for being here. It's my pleasure to introduce two students to you today. John already said quite a few things about this energy lecture series. So it's a student energy lecture series runs in the summer. It's organized mostly by Maxine Lim here in the front. She's run it for about 10 years now. Richard Sassoon, myself, and Steve Aglash from Slack have been the judges this year. It was a pleasure to listen to all the talks and hear the student talks. These four, two this week and two on the 24th, were the ones who rose to the top with the best stories, the best delivery, the most interesting work in our view. But all of them were phenomenal presentations and it was very competitive like John says. So I hope you enjoy these. I find them very interesting. They're just a snapshot of the sort of work that we do in energy or that we fund an energy here at Stanford through faculty and students and postdocs doing the work. So it's my pleasure to introduce Chastity Lee and I will let her tell you about her project. All I will say is that she's a fourth year PhD student here and she's in Mac Cannon's lab and she's doing work turning CO2 into fuels. Thank you so much for that wonderful, very kind introduction and I really appreciate the opportunity to be able to share some of my research with everybody here. And so let's just get into it. First, I want to start off with some really good news which is that there is growing momentum for renewable energy and in particular renewable electricity will continue to grow in the coming years as well as green hydrogen as the price of renewable electricity continues to drop. However, despite increasing electrification, liquid fuels are not going away. In fact, the demand for liquid fuel is projected to increase over time and this is tied closely to the high economic growth. So sectors such as aviation or heavy shipping are not as likely to be able to electrify in the short term. So this really begs the question, can we leverage renewable electricity and green hydrogen in order to regenerate liquid fuels from CO2? And the good news is that there are already existing strategies that convert syngas into renewable liquid fuels. One example here is, you know, you can think of, for example, fissure tropes or CO2 gas fermentation, but in order for this process to be renewable and sustainable, the syngas that feeds these processes needs to come from a renewable source, i.e. from water and CO2. So a straightforward way you may think of to do this is to do a dual electrolysis method to just directly generate CO from CO2 electrolytically and hydrogen similarly from water. However, there are some challenges associated with this process, namely that it has modest efficiency, low current density, poor impurity tolerance, and really no manufactured capability to do this at scale currently. And I wanna contrast that with a combined electrochemical, thermochemical process called a reverse water gas shift. So in this process, we're still generating the hydrogen from water electrolysis, but instead of generating the CO from electrolysis, we're taking some of that hydrogen and CO2 and using a thermochemical process to generate a CO-containing strain. And this is a old reaction, a well-known reaction with well-established reaction engineering and really has an opportunity for quick path to scaling if we can just find a catalyst that meets the requirements necessary. So I've already mentioned how reverse water gas shift can lead us to syngas, but there are some really important things we have to keep in mind. Namely, reverse water gas shift is an equilibrium-limited reaction. So I've plotted that here, the equilibrium conversion of CO2 as a function of temperature, the equilibrium-limited conversion shown there in gray for a one to three CO2 H2 input stream. For the downstream technologies that we're targeting, we're really looking at kind of this temperature regime from 400 to 600 roughly. But actually, under these conditions, the formation of methane through the Sabatier reaction is actually preferred thermodynamically. So the challenge here is to find a catalyst that is 100% selective for CO. I don't need to tell you that, so if you make any methane and you release that, that's gonna negate any possible benefits that you make because methane is so much worse as a greenhouse gas compared to CO2. So for that reason, a lot of the existing work on reverse water gas shift catalysts have focused on the higher temperature region above 600 and so where the methaneation reaction is less favorable. Similarly, they also target less reducing atmospheres, again just to kind of combat that methaneation. So this really gives us a great opportunity if we can find a catalyst that can bridge the following gap, which is if we can take an existing commercial technology that I've already mentioned, water electrolysis and some sort of downstream syngas technology. In this case, for this talk I'm gonna focus on gas fermentation which takes syngas, feeds it to these microbes and then they generate ethanol from that. If we can just find a catalyst that can fill this technology gap that meets these following targets, it needs to be active in the temperature region that we're looking at. It needs to be 100% selective for CO. It needs to be tolerant to many of the impurities that they would be exposed to such as H2S and metal free. If we can do that, we can generate these liquid fuels sustainably. So this is a really great opportunity here and this particular pathway still allows us to maximize renewable to chemical energy conversion but reserves a selective CC bond for biology. And just a quick point of clarification about the gas ratios. From a stoichiometric point of view, any of these gas ratios of CO, CO2 and hydrogen that satisfies this ratio can in theory generate your saturated alcohols. So here I have a list at the bottom here. We have a one to three CO2 to H2 which has the lowest energy content but is the most easily attainable and all the way at the top there we have a one to two CO to H2 which is the hardest to access. So the acetogens for gas fermentation, you can feed them just CO2 and H2 and they'll survive but they won't thrive and they won't produce liquid fuel and they need some CO in order to generate your saturated alcohols but it's not necessary to have pure CO to generate the saturated alcohols. And in fact, the ratio that's targeted is a one to one to five CO, CO2 to H2 ratio and that's the ratio that I'll be targeting for the rest of the talk as well. Our catalyst is a alkaline metal carbonate that is dispersed on a support. We make this catalyst using incipient wetness and prognation where we take one pore volume's worth of a 1.75 molar cesium carbonate solution and very carefully drop dry and mix on a support and a mix and dry. And the idea is that the pores inside the support will be a capillary action pull up that solution and once it dries, the carbonate is dispersed on the surface. We can confirm that it's sort of evenly distributed on the catalyst using energy dispersive X-ray spectroscopy. And additionally, to further characterize our catalyst, we turn to infrared spectroscopy. If you're looking at a bulk carbonate species, the environment around each carbon-oxygen bond is symmetrical so you only get one peak in the IR spectra, there's no splitting. But once that carbonate interacts with the surface, which is what we believe to be the case in our catalyst, that breaks the symmetry and you get a splitting in the peaks and the delta in the wave number can be used to identify which type of binding mode you have. And so our catalyst we see is consistent between somewhere in a bi-dentane and a bridging mode. And we believe that this sort of dispersion of the carbonate on the surface is what leads to the super-basicity of our catalyst. With that in mind, you can perhaps envision a possible catalytic cycle where we have our catalyst here. And as I said, the super-basicity might allow it to deprotonate hydrogen gas forming a metal hydride. The metal hydride can then capture an incoming CO2 molecule forming a formate intermediate, which can then decompose into our product CO. And the formation of water can then regenerate the catalyst and close the loop. To test our catalyst, we had to build a custom flow reactor in our lab. And the basic idea here is that we're flowing in our input gas from the right there. It's going up through down the internal thermocouple into our reactor tube, where there's packed with quartz in our catalyst, and the furnace is surrounding it and is heated to our target temperature. Our product gases will then flow down through the tube and out to the GC for quantification. And any liquid, because we're making water, will be captured in the gravity trap. So for the next series of reactions, the reaction conditions that we targeted are three to one hydrogen to CO2 input stream. We're operating at a pressure of 10 bar and 40 SCC and flow rate, or mills per minute, which with the amount of catalyst that we're loading corresponds to a weight hour of the space velocity of around two. And first, I just wanna show you a control that we did. So in this control, we're stepping through every single temperature step and holding at each temperature for approximately three hours. And as you can see, there's very low reactivity, but there are some because they're stainless steel in the reactor walls, which is reverse water gas shift active. Next, we look at just the bare support. This is titanium with no dispersed carbonate. And we can see there is a slight improvement in the CO yield, but again, we're not nowhere near that equilibrium. And to give you kind of a sense of what the state of the art is, platinum on titanium is one of the best reverse water gas shift catalysts. And as you can see, it does perform pretty well in the low temperature. Doesn't quite make it to equilibrium conversion, but the key problem here, if you look on the right, is that we're generating, as our selectivity is not very high the whole time, we're generating actually a lot of methane, especially at the higher temperatures. And I wanna contrast that with our cesium carbonate on titanium, which even though in the lower temperature doesn't do as well, it is able to reach the equilibrium conversion at around four, maybe 80 or so. And the key point is that it remains 100% selective for CO for the entire duration of the run. And this is a similar story with alumina. It's a different catalyst support material. There's a higher background, but when we load the cesium carbonate on it, we see that it reaches equilibrium conversion in this case at an even lower temperature and still remains 100% selective for CO. You might be curious, what about the stability of this catalyst? So using the same reaction conditions as before, we held these catalysts at a particular temperature in this case, 463 degrees Celsius, and observed that the yield is very stable for the cesium carbonate on alumina. But on platinum, on titanium, we see that there is some decay that might be attributed to the support itself, but the key factor here is that it's still generating a ton of methane over the entire duration of the run. A similar story can be seen at sub-equilibrium temperatures where we're not expecting to reach equilibrium conversion. Again, for 40 plus hours, our catalyst remains stable and the platinum on titanium is producing a ton of methane. So with that, we're able to conclude that our dispersed carbonate catalysts can achieve stable conversion and remain selective for CO even at sub-equilibrium conditions. And also the stability in the presence of an impurity is probed here. So the key difference here is we have still the same three to one H2 to CO, but we're adding 50 ppm of hydrogen sulfide, which is a well-known catalyst poison. And in this case, as you can see, the conversion remains stable for 40 plus hours, even in the presence of H2S. And we hypothesize that the reason this is is because the primarily ionic interaction between the cesium and the sulfur is not stable in the presence of water. And so therefore this prevents us from falling into a thermodynamic sink that can poison the catalyst. We've also looked at, with these same conditions, what happens if we change the identity of the alkaline metal cation. So for reference, here is the cesium carbonate one, which I've shown before. And as we go up the periodic table, we go to potassium and then sodium with decreasing basicity. We see that it follows a trend of lower conversion, but the key point here is that, despite the lower conversion, we're still reaching equilibrium conversion after about 480 or so. We're also looking at different morphologies. In this case, changing to a pellet with an eye towards scaling up. And as you can see, there's slight difference potentially due to the packing of the catalyst, but we're still relatively close to the equilibrium conversion. And I just wanna share that we are able to produce this on a two kilogram batch scale. And if you're curious about some of the efficiencies for water electrolysis, which is a current in commercial technology, the estimated efficiency here is, our energy requirements is 364 kilojoules per mole at H2. For our reverse water gas shift reactor, we're estimating with a really efficient thermal chemical heater that we can have just a requirement of 37 kilojoules per mole of CO. And finally, for this particular downstream application of fermentation plus distillation, with the current technology, the energy requirements are around 370 kilojoules per mole of ethanol. And if you do all the math, do all the stoichiometric conversions, and then the efficiency, that leads us to 2.8 megajoules per mole of ethanol, which if you take the higher heating value for ethanol is roughly 50% power to ethanol efficiency. And so I just wanna really quickly share some of my future work. So we are hoping to, are actually in the process of building a much larger one inch diameter tube here at Stanford to be able to evaluate the morphologies of those more industrially relevant catalysts. And we also are going to be moving forward with testing with high throughput against industrial catalysts, all towards an eye towards de-risking these technologies for commercial incorporation. And with that, I really just would like to thank my PI Matt, who has just been really inspirational and helpful this whole time. Dr. Amy Frankhauser, who did a lot of the pioneering work in this, and Keisha, who has joined me on this project and is doing a lot of great work already. And obviously my two funding sources, the Chevron Fellowship and the Stanford Sustainability Accelerator. And of course, everyone else at the Canon Lab. So with that, I'll be happy to take any questions. Thank you. Lovely job, thank you Chastity. So we're probably gonna hold the questions until the end, if that's okay. That will next invite Eric McShane up. So I'll introduce you if you wanna start getting your presentation ready. So Eric is actually a postdoc here. He's in Mateo Curgnello's lab, and he's working our project to produce fertilizer in an alternative way to what we currently do now, which is using the Haber-Borsch process. So he's gonna tell you all about that. He's recently come here, I think you graduated in September 2021. So I don't know exactly when you arrive, but it's a fairly recent arrival. So Eric, on you go, take it away. Thank you, Jenny, for the kind introduction. And I'm excited to present today my work on trying to electrify the traditional Haber-Borsch process. My title of my talk today is quantification of salt electrolyte interface composition during non-aqueous electrochemical nitrogen reduction. And so to briefly outline my talk, I'll first get into why we really care about nitrogen reduction in the first place, why it might make sense to electrify the traditional process of nitrogen reduction, and then delve into specifically one constituent of this process, the solid electrolyte interface that is formed in such reactions. And I'll also quantify specific species that are formed in this solid electrolyte interface layer. And finally, I'll conclude with some remarks and future directions for this project. So to begin, first of all, nitrogen reduction has been an extremely pivotal reaction, actually, in the evolution of the human species, I would even argue. In particular, the Haber-Borsch process has been attributed with really feeding the world population over the last century or so since it's existed. And this specifically turns dinitrogen and hydrogen into ammonia and H3. Shown here is some estimates of what the world population may have looked like in the red curve without the advent of the Haber-Borsch process. So it would approximately double, say, in the last century or so. But instead, what we actually have observed in blue is the approximate quadrupling of the world population in large part thanks to the Haber-Borsch process, which provides synthetic fertilizer to feed most of the world. And then therefore, the gray is what we can specifically attribute to the advent of the Haber-Borsch process. But what I have neglected to mention thus far is that this reaction is not necessarily the easiest reaction to complete. It specifically requires moderate temperatures, say about 400 degrees Celsius, as well as particularly high pressures in the range of about 150 bar or so. So these demanding conditions, therefore, result in the necessity of large centralized plants to conduct the reaction. It also requires a large amount of energy. By most estimates, around 1% of global energy consumption can be attributed to the Haber-Borsch process. And an approximately commensurate amount of CO2 emissions are therefore attributed, likewise to the Haber-Borsch process as well. And to delve a little bit deeper into the implications and reasons why we might care about how this population is growing, I'll specifically show up here the global population density. And the first thing that we might notice is that population hubs tend to be around the coastal regions of a given continent. And in particular, the exceptions to this are the entirety of India and the entirety of China are fairly densely populated. But one particular region I'd like to draw your attention to is the region in the red box, the Sub-Saharan Africa region. This is notably a more distributed population such that the population centers span a very large geographic region throughout the entirety of Africa. And it becomes uniquely challenged in some aspects because of this. First of all, the population in this region has grown quite dramatically in the last decade or so. It's grown about 40% in population. And there is unfortunately not a very reliable form of energy generation throughout the majority of this region. But that does give us an opportunity then to engineer perhaps cheap, modular, meaning that it can be scaled up or down quite easily, forms of energy generation and processes that are amenable to this type of energy generation. And so, taking this all into account, we want to use process electrification which enables these modular designs, for example, ammonia synthesis, to specifically address the needs of a region such as Sub-Saharan Africa which has a very dispersed population. And so then we'll get into how and why we might electrify the nitrogen reduction process. First, I'd like to point out that in addition to the obvious environmental benefits of using renewable energy, there is increasingly becoming an economic incentive to use such processes of generating energy. Shown here is the solar PV cost over the last decade or so. And it's showing a logarithmic or exponential decline in cost over the last decade. You can actually extend this to the preceding decades as well and it follows a very similar trend. And therefore we expect it to perhaps continue a similar trend in the future. In addition, wind turbines have similarly benefited from exponential cost decline, especially in the recent decades, as have storage mechanisms such as lithium-ion batteries used for grid-scale storage as well. So putting all this together, actually right now, if you're to calculate the levelized cost of energy which is simply the cost per unit of energy produced, we find that solar and winds are the cheapest form of generating energy right now. However, what I will say is this graph does not necessarily account for the additional cost of lithium-ion battery storage to sort of even out the intermittency of such generation methods. And if you're to include lithium-ion batteries as well, in some regions it may be cost-effective, however in some regions it's more expensive than natural gas, for example. But if we're to project these cost declines into the future, we can see a future where indeed renewable energy makes the most sense economically and environmentally. And so our goal then is to leverage these existing trends and develop an electrified decentralized manner of producing ammonia. And I'll get a little bit deeper now into the exact method we use to actually produce this ammonia electrochemically. First, I'll show the four main regions of such a system. The oval region is the electrolyte solution. There's then some sort of interfacial layer that exists on the cathode and also a nitradation layer, which is shown in pink and a dark gray, which is simply the cathode surface on which the reaction occurs. So in the electrolyte, of course, we need our reactant, nitrogen gas. This is in general bubbled through an electrolyte solution. In addition, this electrolyte is predominantly comprised of a THF or tetrahedra furane solvent in 99 volume percent. And this electrolyte additionally contains a lithium salt, particularly people often use lithium perchlorate in about a 0.5 molar concentration range. And because we have this lithium salt when you apply reducing potential, you can deposit lithium metal on your cathode surface. In the presence of nitrogen, this can form a lithium nitride type layer on your reactive surface. And we additionally add ethanol into our electrolyte to provide a source of protons, through which we can add protons to this nitride layer. And if we do this enough times, first of all, we sequentially form ethoxide, the anionic form of ethanol. And eventually if we do it three times, we can form our desired product, the ammonia molecule. And what I have neglected to mention so far though, is that this reaction is not 100% selective towards our desired product ammonia, rather some byproducts can be formed. And these form the basis for the so-called solid electrolyte interface, or SEI layer that forms on the cathode. And this will be the subject of my studies that I'll present. So delving a deeper into this SEI layer, specifically in the electrochemical nitrogen reduction reaction. If we use the traditional electrolyte composition, which employs one volume percent ethanol, we can imagine what might happen to this ethanol species if it is directly reduced on our cathode surface, as opposed to donating its proton to form an eventual ammonia. If ethanol is directly produced, we imagine that hydrogen gas will be evolved, and you will in turn form an ethoxide species that is anionic. And therefore, this anionic species must find some sort of cation to counterbalance its charge. One option is it could be protonated by a proton that may exist in solution. This would form ethanol, which we know is a liquid, and would therefore be a constituent of our electrolyte. And also, one possibility is that it can find a lithium counterion, and we know that lithium ethoxide at room temperature is a solid, and could thus precipitate on the cathode surface and form the basis for this SEI layer, and we'll delve into that a bit deeper going forward. In addition, THF is present in 99 volume percent in the electrolyte solution, and we can imagine that if THF is directly reduced on our cathode surface, it can undergo a reductive ring-opening reaction where radical butoxide species is formed. This can undergo hydrogen abstraction, followed by perhaps a protonation or lithiation similar to ethoxide. And similarly, if the protonation occurs, we form a liquid species, butanol in this case. Or if it's lithiated, we form another solid species that might comprise a portion of our solid electrolyte interface layer, lithium butoxide. So with all this in mind, my goal and my work was to quantify these SEI species with the aim of understanding degradation or byproduct mechanisms, as well as learning the effect of this SEI on the ultimate selectivity of our reaction. In order to do this, I used this shell shown here, and I'll first point out that a bubbler is stuck into our cell. In general, we bubble nitrogen gas through our bubbler to supply a steady stream of nitrogen to be reduced to ultimately form ammonia. But in addition in my systems, I also flow hydrogen gas with the ultimate goal of oxidizing the hydrogen gas to additionally supply a steady stream of protons that can be used to protonate the nitrides formed in our system. In addition, I have a copper cathode on which the nitrogation layer forms and ammonia is ultimately produced. And this is all submerged in about 16 milliliters of electrolyte in my particular system. Here shown is just an optical image of the top portion of the cell that I use with the cathode shown clearly. And in general, I use a platinum mesh anode on which we hope hydrogen oxidation predominantly occurs. And a reference electrode, just for note, it is a half lithiated iron phosphate which allows us to separately probe the cathode and anode during the reaction. And when we run such a reaction, we can extract the cathode out of the cell at the end of a reaction. We can put it in a septum cap vial that is loaded in an argon glovebox such that it remains air free. And we can take that septum cap vial out of the glovebox, inject D2O into the vessel, and then analyze the rinsates with NMR. The theory being that any SEI species that were formed will be rinsed off by the D2O rinse. And we can ultimately quantify what ends up in the rinsate with an NMR technique. And so with all this in mind, I turn to how this cycling data actually works and how we extract at varying time points throughout a reaction to get the data that we're interested in. So here shown is just the simple voltage profile of our two electrodes versus the lithium iron phosphate reference. Noting the conditions include the aforementioned electrolyte composition, the aforementioned cathode anode and reference electrode, and a flow rate of five mil per minute of 5% hydrogen and nitrogen. Once again, the hydrogen is added so that we can oxidize hydrogen to provide a steady stream of protons throughout the reaction. In addition, we apply a 1 milliamp per centimeter squared current throughout these reactions in order to delve into the effect of varying amounts of charge pass on our existing SEI species. So first, the blue and the red represent the anode and cathode potentials respectively. We note that at short times, we observe around a one volt versus our reference potential for the anode and around a negative four volt versus our reference for the cathode. If we let this reaction proceed even further, we find that the anode potential remains relatively stable around one volts. However, the cathode potential at say around two to three hours appears to systematically drop to around negative six volts or so. And if we let this reaction proceed even further, the anode remains at around one volt admittedly with some noise, likely due to some bubbling effects. And the cathode seems to level off at around negative five volts or so. Proceeding all the way to 24 hours, we find that this in general continues throughout the duration of such experiments. And I'll note that we collect our sample in these different cells after different amounts of time and therefore amounts of charge pass. And we analyze the rinse rate of NMR to determine what SEI had formed in each of these different experiments. And so first, just show you some data for the two hour mark, which corresponds to two milliamp hour of total charge pass. In general, NMR data is plotted as intensity versus chemical shift. And I'll note that chemical shift for our purposes just corresponds to a characteristic response of different species in solution in response to an external magnetic field. And therefore by determining where the chemical shift lies for a given species, you can determine what the identity of a species is. And the intensity in general for our purposes just related to the quantity of a given species in solution. And therefore we can use such a technique to quantify various rinse rate species. Here shown is the NMR spectra collected after the two hour mark in our experiments. There are two prominent peaks I'd like to describe first. One is the water peak, which is inevitably present in a D2O rinse due to impurities in the D2O. And I intentionally added a DMSO internal standard with which I can describe the position of other peaks relative to. I'll note that the largest non-standard peak that we observe is an ethoxide related peak. And I've listed it here as deuterated ethanol. Presumably this species was lithium ethoxide on our cathode surface. And then when we rinse it off, it may become either the deuterated form of ethoxide or perhaps it's still lithium ethoxide that we can likewise still measure with NMR. And there are many other smaller peaks that are barely visible at this scale. But I'll note one for reference, this is likely a propoxide related species, which means that we likely formed lithium propoxide in our reactions, which was not necessarily anticipated but we can nonetheless see with NMR. And there are other small, minor species that are present in our SEI layer and therefore I'll zoom in on this particular region to delve a bit deeper into that. Here shown is that aforementioned zoomed in view of the previous slide. We note that the ethoxide large peak is now largely off of our screen. And we do see some satellite peaks for the ethoxide. This can occur, for example, if you have a large amount of ligament species and there are different natural abundances of de-duration of that given species and solution. In addition, we see the propoxide species that we mentioned on the previous slide, but also at this magnification, we can indeed confirm the presence of a butoxide, which we anticipated from our THF reduction process, as well as some very small peaks that we can likely, in this case, ascribe to, for example, an isobutoxide. But like I said, some of these are barely perceptible at our magnification and conditions shown here. But nonetheless, this gives us a little glimpse into what other minor SEI species may be forming, and therefore helps us learn about mechanistic degradation or side products formed in such a reaction. And so if we're to look at how this changes over the course of a 24-hour cycling experiment, we see that at about 0.5 hours, we don't see much in our NMR spectrum besides the internal standard. But at 24 hours, there seems to be a monotonic increase in the amount of ethoxide even viewed at this magnitude of the NMR spectra. And what's more, we can look at some of these minor SEI species by zooming much farther in. And this will be shown on the right-hand side of your screen. We can see that in general, there are a lot of minor species that we can observe with NMR. This is great to further our mechanistic understanding of what sort of side reactions may be occurring. But I have not so far exhaustively gone through and identified each species. But for our purposes, we can at least see that all of these minor species seem to be growing as the reaction proceeds. And so, in general, we can take these NMR spectra and we want to be able to quantify each of these species that are present in solution. On this next slide, I'll show that we can easily quantify because the peak is so apparent and large, the ethoxide that is formed in each of our experiments. And when we do so using an external standard of ethanol and water, we find that the ethoxide growth seems to follow a trend shown here where there is a tapering off in the growth after about 10 hours of reaction time. And the first takeaway that we get is that, say, after about half an hour, there's no apparent ethoxide that's been formed. This likely arises because perchlorates, which may be present, is maybe reduced to form a chloride species and therefore no reaction involving ethanol is actually occurring in the first 30 minutes of reaction time. There is, however, rapid growth of the amount of ethoxide present between the 0.5 and 10 hour mark. And as we just alluded to, this seems to wane after the 10 hour mark, indicating perhaps that something is preventing continual ethanol reduction. And with this in mind, I've now plotted that same data on the previous slide, but just a larger y-axis scale to capture some other trends that we have not yet discussed. And if I plot, in addition to the charge pass that goes to ethoxide, the ammonia that is created, the charge pass that goes to ammonia specifically, we find that although the ammonia production rate seems to be modest for the first, say, three hours or so, it really begins to take off, say, around three to 10 hours, especially at 24 hours. And if we're to show specifically what's happening at 24 hours, we know that at a 1 milliamprecentimeter squared rate, we pass 24 milliamp hour worth of charge and 6 milliamp hours of that charge goes directly towards producing ammonia. And so we can then define what's called a Faraday efficiency, which is simply the percentage of the amount of charge pass that goes towards our desired product, which is in this case ammonia, of course. When we plot this for the various time points, we find that especially at short times, the Faraday efficiency is quite modest around 10%, but this really seems to take off later in the cycling and it seems to plateau between the 20 and 25% Faraday efficiency mark at long times. And this is quite interesting because this seems to hint at the role of this ethoxide layer. Specifically, what we might be able to take away here is that the ethoxide layer, although its growth does not continue all throughout the reaction, it may seem to be acting sort of like a membrane such that the ethoxide layer is selectively impeding further ethanol transport to the surface, meaning that the side product of direct ethanol reduction does not occur as substantially later in cycling, but is still allowing the smaller molecule of dinitrogen to reach the surface to undergo the desired nitrogen reduction reaction. And so with all this in mind, I'd like to just conclude and discuss what we've learned so far and maybe some future directions for this project going forward. First of all, I'd like to say that lithium ethoxide we found in our work is indeed the major SEI component as a result of the ethanol reduction reaction. This growth of the ethoxide layer seems to take off between half an hour and 10 hours and seems to really wane in the remaining 14 hours of reaction time. And in general, the selectivity towards ammonia seems to increase at later times in the reaction, suggesting perhaps that this ethoxide layer might be beneficial towards steering selectivity towards ammonia in particular. And with that in mind, some future directions are interesting to pursue as well. We can quantify, of course, the minor SEI species which we can observe but cannot yet quantify due to lack of external standards so far, but we fully intend to do that. We can also examine the effect of cycling parameters using this novel characterization technique of the SEI layers. And we can delve into how alternative proton donors, which have previously been found, to really alter the selectivity of ammonia production may in fact influence the SEI layer as well. And last but not least, I'd like to just tie this back in once for all to the big picture. Haber-Bosch has been a pivotal invention over the last century, but over perhaps the last three decades or so, there has been a leveling off of the performance improvement, specifically the energy efficiency that we attribute to the Haber-Bosch process. Approximately, if we calculated out, it's about 60% energy efficient at the current moment, and it doesn't seem to be drastically improving in the recent years. Right now, one opportunity to get past this threshold is to use a lithium mediated or electrochemical approach, which we've shown. And right now, the state of the art, especially in recently shown papers, is around, say, 25% energy efficiency. This is obviously less than the traditional Haber-Bosch process, but it has some distinct advantages over traditional Haber-Bosch. First of all, this is amenable to small distributed plants because it does not require extremely high pressures to operate and can therefore be scaled down easily. And it also uses seamless integration with what we anticipate will be the major form of energy generation, such as renewable solar and wind. And lastly, because this is not yet at parity with traditional Haber-Bosch and it's rapidly improving in energy efficiency, this represents really a vast white space where academia in particular can add extreme value to this emerging technology because it's not yet commercially applicable. However, it is improving rapidly and may eventually surpass even in terms of energy efficiency, traditional Haber-Bosch. So with all this in mind, I'd just like to acknowledge a few key people. First of all, Matteo, who's in the audience today, has been instrumental in allowing me to really delve into a topic deeply, has helped my evolution as a scientist, has really allowed me to just pursue my passions in lab. In addition, I'd like to thank my group members who have been also really helpful in discussions, whether at lunch or in the office or even in group meeting or subgroup meeting, as well as some specific people in the Haramu Lab who have provided support on the project as well. And finally, my funding source, the Villain Foundation, for their financial support on this project. Thank you all for attending. I'd be happy to take any questions. Lovely job. Thank you, Eric. And thank you, Chastity, again. So any questions? I don't know how much time we have, but we'll keep going until we're told otherwise, so. Just working. From what I understand, the natural gas feedstocks right now are most feedstocks for natural gas production, mostly natural gas. Have your processes the same feedstocks? So feedstocks, one can think of feedstocks to generate hydrogen or to provide energy. And I would say actually both would involve natural gas. Yes, currently they involve natural gas. So the inputs are the same, but just the process itself is what your research is. Right, right. So this process still, as of now, requires hydrogen. There are potential ways to eliminate that need in the future. The hope, one way, we could provide hydrogen is like electrolyzing water, for example, which would significantly reduce any sorts of emissions related to a fossil fuel-based process. Yeah, but as of now, I'm not exploring actively hydrogen production methods for this process in particular. And I guess a quick second question. What do you mean by seamless integration with renewable energy? What's different? Right, right. So one can imagine that existing Haber-Bosch plants can have a heating method, for example, by burning a fossil fuel to provide the energy. However, this process is inherently electrochemical in nature, and therefore it requires using electrons, and therefore you can seamlessly integrate any type of electrochemical process with renewable energy, which forms electrons as the charge carrier. A couple of questions for Chastity, really nice work. The first one is, do you have evidence that the CO that you're producing comes from CO2 reduction, and not from decomposition of your carbonate? So this is a good question. And I think, so we're in a flowing CO2 atmosphere. So, and these runs, they go for multiple days, right? So if it were purely from decomposition, which I think there probably is some amount of it, even if there is, I think there would be CO2 kind of taken back in and reforming the carbonate. We also see the generation of a formate intermediate sometimes when we take the NMR of the rinse afterwards. So we do, for that reason, we think it probably is due to the CO2. And I think, sorry, to just kind of, for, we can run this for upwards of a week and we still see the same conversion over time. Whereas if it is coming from the carbonate decomposition, we'd expect that to kind of tail off, right? As long as you, if you have a lot of carbonate, then you would also make a lot of CO. So if you were able to label the CO2, for example, it would be able to. Yes, that would be great. Or you run it with just hydrogen, and then if you do see CO, then you know that it's coming from your carbonate. But I do believe that you make CO. Then the other question is more on the mechanistic aspect. Splitting hydrogen is very hard. Yeah. Because hydrogen has a strong bond. It's not very acidic. Yeah, and you are proposing that hydrogen splitting is part of your mechanism. Do you have evidence for that? Yes, so actually, I wish I had my slide, but the sort of, if going back to kind of what led us to this technology in the first place. So we started out with this catalyst reacting with benzene, which the gas phase acidity of benzene is very, very similar to hydrogen. I think it's like 40 and 41, I forget which one is which. But they're very similar. And we were able to show that these types of catalysts, these dispersed carbonate catalysts are able to deprotonate benzene. And therefore we think it's pretty, because the gas phase acidity of hydrogen is very similar, we think it's also likely that it's also able to deprotonate hydrogen. I think if you run it with deuterium, and you observe a kinetic isotope effect, then you would know it's really hydrogen splitting is the limiting side, for example. Yeah, that would definitely confirm that. Yes, so yeah, if I get the chance to, I would love to run that. Time for one more. Great presentations, both of you. Question for Eric. Have you taken one of these copper cathodes with the SEI generated already and run it again after to see if you don't see this passivation happening since it's already there? Yeah, I have not done that, but that's a good suggestion. It's even come up in our subgroup meetings, but yes, I think that would be quite interesting to see if it survives, for example, the rinsing process and being put back in a cell. That would be quite interesting. That means you really only need to form it once to form this SEI layer, and then it should persist indefinitely. Yeah, that's a great suggestion. Associated with that, do you have an idea of the mechanism of exclusion of ethanol by this lithium ethoxide membrane? Good question as well. So in general, dinitrogen is a smaller molecule than ethanol, and therefore we think it's some sort of size selective, poor size effect of this SEI layer. Yeah, you can imagine that this SEI layer is like a bunch of perhaps spheres that are packed in some given porosity. If the poor size is such that nitrogen can traverse it, but ethanol can't, that's what I'm kind of imagining for this selective membrane type behavior. All right, thank you. Yeah, yeah. Wonderful, thank you very much for your questions, and I'd ask you to thank your speakers one last time. And then I am, thank you. Their prize for being some of the best speakers, student speakers, is to present here to you guys, but it's also to receive this nice little award plaque that our Nancy Sandoval prepared for them. So thank you all very much. As I do this, you're free to leave. Thank you.