Thermodynamic Assessment of Microencapsulated Sodium Carbonate Slurry for Carbon Capture

Available online at www.sciencedirect.com

ScienceDirect Energy Procedia 63 (2014) 2331 – 2335

GHGT-12

Thermodynamic assessment of microencapsulated sodium carbonate slurry for carbon capture Joshuah K. Stolaroff1 * and William L. Bourcier1 1

Lawrence Livermore National Laboratory, Livermore, CA 94551, USA

Abstract

Micro-encapsulated Carbon Sorbents (MECS) are a new class of carbon capture materials consisting of a CO2absorbing liquid solvent contained within solid, CO2-permeable, polymer shells. MECS enhance the rate of CO2 absorption for solvents with slow kinetics and prevent solid precipitates from scaling and fouling equipment, two factors that have previously limited the use of sodium carbonate solution for carbon capture. Here, we examine the thermodynamics of sodium carbonate slurries for carbon capture. We model the vapour-liquid-solid equilibria of sodium carbonate and find several features that can contribute to an energy-efficient capture process: very high CO2 pressures in stripping conditions, relatively low water vapour pressures in stripping conditions, and good swing capacity. The potential energy savings compared with an MEA system are discussed. © 2013 2014 The by by Elsevier Ltd. Ltd. This is an open access article under the CC BY-NC-ND license © TheAuthors. Authors.Published Published Elsevier (http://creativecommons.org/licenses/by-nc-nd/3.0/). Selection and peer-review under responsibility of GHGT. Peer-review under responsibility of the Organizing Committee of GHGT-12

Keywords: carbon capture; CCS; CO2 ; sodium carbonate; potassium carbonate; MECS; microcapsules; micro-encapsulation

1. Introduction Micro-encapsulated Carbon Sorbents (MECS) are a new class of carbon capture materials consisting of a CO 2absorbing liquid solvent contained within solid, CO2-permeable, polymer shells. The resulting capsules are spherical and highly uniform in size. They have been produced with diameters of 100—600 m using microfluidic devices,

* Corresponding author. Tel.:+1-925-422-0957 E-mail address: [email protected]

1876-6102 © 2014 The Authors. Published by Elsevier Ltd. This is an open access article under the CC BY-NC-ND license (http://creativecommons.org/licenses/by-nc-nd/3.0/). Peer-review under responsibility of the Organizing Committee of GHGT-12 doi:10.1016/j.egypro.2014.11.253

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which afford excellent control over MECS materials and geometry. MECS combine advantages of solid sorbents (high surface area, low volatility) and liquid solvents (high capacity, water tolerance, thermal regeneration). MECS enable the use of CO2 solvents that are otherwise impractical in conventional reactors, such as those with high viscosity or corrosivity, slow absorption kinetics, and those that precipitate solids [1—2]. Sodium carbonate solution is an excellent candidate for encapsulation. It has several desirable features for use in carbon capture, such as low cost, zero volatility, high capacity and favourable thermodynamics [3]. However, use of sodium carbonate has been limited by its low solubility (tendency to form precipitates) and slow absorption rate of CO2. Encapsulation mitigates both these drawbacks. MECS containing sodium carbonate solutions have been previously produced and shown to have a ~10 times higher CO 2 absorption rate compared with sodium carbonate solutions in a typical liquid stripping tower, due primarily to increased surface area. These MECS have also been shown to maintain integrity and capacity over multiple thermal regeneration cycles. [2] Here we assess the thermodynamic properties of a capture system based on sodium carbonate solution. In particular, we examine the vapour-liquid-solid equilibrium (VLSE) of sodium carbonate slurries at relevant concentrations and temperatures to a precipitating capture process. Although potassium carbonate has been examined in some detail because of its use in the Benfield process [4], sodium carbonate has so far received less attention because of its lower solubility. Knuutila et al. [5] reviewed available data sources and measured vapour pressures of the system, but did not find or measure concentrations high enough to include solid phases. As we will show, the VLSE behaviour with solids present differs from the lower concentration behaviour in ways important to carbon capture. In this paper, we describe our methods for calculating the VLSE and choice of input parameters. We then take a solution of concentrated (30 wt%) sodium carbonate through a representative carbon capture cycle, from CO2 absorption to heating, CO2 desorption, and cooling, with reasonable loss and uptake of water. Our parameter choices are informed by the concept of MECS in either a fixed bed or staged, fluidized bed with high-pressure stripping. However, the results represent chemical equilibria that are not specific to MECS and also would apply to, for example, a spray tower configuration. 2. Methods We model the VLSE of sodium carbonate slurries using the Geochemist’s Workbench (GWB) software and the thermo_da0ypfR2 database, which is a Pitzer-type solution model including available high temperature data. The Pitzer model of ion activities allows accurate predictions at much higher ionic strengths than other models such as those based on simpler forms of Debye-Huckel theory. The solubility products for Nahcolite and Trona in the data file were adjusted slightly to better fit carbonate solubility data [6]. The modeling was carried out by first equilibrating a known concentration of carbonate solution with a partial pressure of carbon dioxide of 0.1 bars and 40 oC, representing the carbon dioxide pressure of coal plant flue gas. The loaded (fat) solution was then heated to the stripping temperature and allowed to de-gas carbon dioxide and water in amounts consistent with their partial pressures at the stripping temperature. The de-gassed (lean) solution was then cooled back to 40 oC, and water restored to the original mass to begin a new cycle. The GWB model does not account for a separate gas phase, but does provide partial pressures of gases in equilibrium with the fluid phase. We have assumed equilibrium between the aqueous and gas phases in our model but note there will be so me disequilibrium in a practical process due to kinetic limitations. In the following plots and discussion, “CO2 loading” is defined as the percentage of stoichiometric capacity, also known as the “percent conversion” of carbonate to bicarbonate.

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3. Results We start with a solution of 4 molal Na2CO3 (30 wt%) and expose it to CO2 gas of increasing partial pressure at 40 C. The relationship between CO2 loading and partial pressure is shown in Figure 1. We find that the solution reaches equilibrium with 0.1 bar of CO2 at 93% loading. This is roughly equivalent to an 80% approach to equilibrium with a typical coal fired flue gas (12-13% CO2 at atmospheric pressure). We take this to be the loaded, or “fat” slurry. This assumes isothermal absorption, though in practice the solution will change temperature as it absorbs CO2. The added heat in the absorber is significant for amine solutions, requiring intercooling in some designs. However, the heat of reaction is about half as much for liquid carbonate solutions. And as solid Nacholite forms, an endothermic reaction that begins at about 23% loading (note the discontinuity in the pressure curve) the sensible heat added to the solution is further reduced. Overall, we calculate a temperature change on the order of 5 C in adiabatic conditions. High purity CO2 is removed from the slurry by heating. The pressure of CO 2 and ratio of CO2 to water vapor depends on the stripping conditions. We can consider two endpoints. (1) The slurry is heated at constant pressure and the evolved gas is sent to the condenser and compressor. We assume that each parcel of evolved gas is at equilibrium with the slurry at a particular temperature. (2) The slurry is heated in a fixed volume and then “flashed” gradually to lower and lower pressures. We assume the parcels of gas evolved are at equilibrium with the solution at a particular loading. The temperature remains constant by addition of sensible heat as needed. This case is an idealized multi-pressure stripper. Figure 2 shows the first scenario and Figure 3 shows the second. Note that the CO2 pressure in both cases goes very high, much higher than most CO2 solvents at their respective operating conditions. It appears that the presence of the solid drives this trend. Tosh et al. [4] provide similar data for potassium carbonate (which does not form precipitates) showing CO2 pressure rising more to the range of 10 bar, rather than 100 bar, for comparable temperatures and loadings. Also note the discontinuity in the slope in Figure 2 at 139 C, where the solid dissolves. The slope of the P-T curve turns more gradual once the solid dissolves, becoming similar to the slope for potassium or for low-concentration sodium.

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CO2 loading [%] Figure 3: Removing CO2 from 30 wt% MEA solution at 120 C. Compare with previous figure for carbonates. Data from Aronu et al. [7].

Using Figure 3, we can find the composition of the unloaded, or “lean” solution, which is determined by how high we heat the slurry and which is the lowest pressure that we extract the evolved gases. In practice, these are design choices that would be selected in an optimization. For our purposes, we choose 140 C and 1 bar CO2 partial pressure (about 4 bar total pressure). This yields a lean solution with 38% loading and a swing capacity of 9 wt% CO2. In a capture system, we generally want to minimize the water vapor sent to the condenser, because a significant portion of the energy spent in the stripper goes to vaporize that water. By taking the integral of the pressure curves for CO2 and water, we can estimate the composition of the gas sent to the condenser. For the interval between the rich and lean loadings (the whole of Figure 3), we find a CO2:H2O ratio of about 10:1. Figure 4 shows a comparable pressure vs. loading curve for monoethanolamine (MEA) in stripper conditions. The same integration here gives a CO2:H2O ratio of about 2:1. This suggests we can expect ~5-fold reduction in parasitic water evaporation in a sodium carbonate system compared with MEA. The water lost in the stripper would change the solution composition as it evaporates, however, we find this is a minor effect. Applying the 10:1 ratio, we would expect to lose 0.26 M of water, barely changing the solution concentrations. The quantity of water exchanged with the humid flue gas in the absorber is similarly small. 4. Discussion We have assessed the vapour-liquid-solid equilibrium characteristics of sodium carbonate slurries for carbon capture. They appear to have several features that would enable an energy-efficient process compared to amine solutions: very high CO2 pressure at high temperature (coupled with high temperature stability), comparatively low water pressures at those temperatures, and lower heats of reaction. While an integrated process analysis is needed to estimate the energy use of the complete capture system, we can make some instructive comparisons with an MEA system. Oexmann et al. [8] provide a useful breakdown of parasitic energy load in an optimized MEA capture system. Compression contributes 25% to the total load, water evaporation for stripping contributes about 22%, and the heat of reaction about 18%. Suppose we strip CO 2 from a sodium carbonate system at an average of 10 bar, then we would save roughly 2/3 of the compression energy compared with compression from an MEA stripper at atmospheric pressure. As discussed above, we would also avoid the vast majority of water evaporation. We should save also on the heat of reaction, though the amount isn’t fully known from this analysis. On the other hand, we would expect a higher load for sensible heat transfer (about 27% of total) than MEA because of higher temperature stripping and because the slurry form of sodium carbonate, either as MECS or in some other process configuration, could lead to lower-efficiency heat exchange between the

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absorber and stripper. If we assume the increase in sensible heat is only proportional to the higher stripping temperature (140 instead of 120 C), then we would expect an overall energy savings around 36% compared with MEA. This estimate is consistent with the earlier finding from Knuutila et al. [3] that a sodium carbonate process could operate on a coal power plant with a 9% efficiency penalty, compared with 12—15% for an MEA process. Of course, this very rough calculation leaves out many details of the process design that must be included in the next stage of analysis. However, the VLSE data suggest significant potential for energy savings using the sodium carbonate system along with the inherent advantages in material cost, environmental friendliness, and scalability. 5. Acknowledgements This work was performed under the auspices of the U.S. Department of Energy by Lawrence Livermore National Laboratory under Contract W-7405-Eng-48. This document was prepared as an account of work sponsored by an agency of the United States government. Neither the United States government nor Lawrence Livermore National Security, LLC, nor any of their employees makes any warranty, expressed or implied, or assumes any legal liability or responsibility for the accuracy, completeness, or usefulness of any information, apparatus, product, or process disclosed, or represents that its use would not infringe privately owned rights. Reference herein to any specific commercial product, process, or service by trade name, trademark, manufacturer, or otherwise does not necessarily constitute or imply its endorsement, recommendation, or favoring by the United States government or Lawrence Livermore National Security, LLC. The views and opinions of authors expressed herein do not necessarily state or reflect those of the United States government or Lawrence Livermore National Security, LLC, and shall not be used for advertising or product endorsement purposes. 6. References [1] Aines, R. D.; Spaddaccini, C. M.; Duoss, E. B.; Stolaroff, J. K.; Vericella, J.; Lewis, J. A. & Farthing, G. "Encapsulated Solvents for Carbon Dioxide Capture ." Energy Procedia, 2013, 37, 219 – 224. [2] Vericella, J. J.; Duoss, E. B.; Stolaroff, J. K.; Baker, S. E.; Hardin, J. O.; Lewicki, J.; Glogowski, E.; Floyd, W. C.; Valdez, C. A.; Smith, W. L.; Jr., J. H. S.; Bourcier, W. L.; Spadaccini, C. M.; Lewis, J. A. & Aines, R. D. "Encapsulated liquid sorbents for carbon dioxide capture." Nature Communications (in review). 2014. [3] Knuutila, H.; Svendsen, H. F. & Anttila, M. "CO2 capture from coal-fired power plants based on sodium carbonate slurry; a systems feasibility and sensitivity study." International Journal of Greenhouse Gas Control, 2009, 3, 143-151. [4] Tosh, J. S.; Field, J. H.; Benson, H. E. & Haynes, W. P. "Equilibrium study of the system potassium carbonate, potassium bicarbonate, carbon dioxide, and water." US Department of the Interior, Bureau of Mines, 1959. [5] Knuutila, H.; Hessen, E. T.; Kim, I.; Haug-Warberg, T. & Svendsen, H. F. "Vapor-liquid equilibrium in the sodium carbonate-sodium bicarbonate-water-CO2-system." Chemical Engineering Science, 2010, 65, 2218 – 2226. [6] Monnin, C. and Schott, J “Determination of the solubility products of sodium carbonate minerals and an application to Trona deposition in Lake Magadi (Kenya)” Geochim. Cosmochim. Acta, 1984, 48, 571-581. [7] Aronu, U. E.; Gondal, S.; Hessen, E. T.; Haug-Warberg, T.; Hartono, A.; Hoff, K. A. & Svendsen, H. F. "Solubility of CO2 in 15, 30, 45 and 60 mass% MEA from 40 to 120C and model representation using the extended UNIQUAC framework." Chemical Engineering Science, 2011, 66, 6393 – 640. [8] Oexmann, J.; Kather, A.; Linnenberg, S. & Liebenthal, U. "Post-combustion CO2 capture: chemical absorption processes in coal-fired steam power plants." Greenhouse Gases: Science and Technology, John Wiley & Sons, Ltd., 2012, 2, 80-98