Molten Salt Promoting Effect in Double Salt CO2 Absorbents - The

Dec 21, 2015 - Lixia YangDan LiuPingping WangHwimin SeoJianzhou GuiYong-Ki Park. Industrial & Engineering Chemistry Research 2018 57 (31), 10591- ...
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Molten Salt Promoting Effect in Double Salt CO2 Absorbents Keling Zhang,†,§ Xiaohong S. Li,‡ Haobo Chen,‡,⊥ Prabhakar Singh,† and David L. King*,‡ †

Center for Clean Energy Engineering and Materials Science and Engineering, University of Connecticut, Storrs, Connecticut 06269, United States ‡ Institute for Integrated Catalysis, Pacific Northwest National Laboratory, P.O. Box 999, Richland, Washington 99354, United States ABSTRACT: NaNO3 and other alkali nitrate salts, which are present in the molten state during use, have been described as facilitators or catalysts for CO2 absorption by both MgO and MgO-containing double salts. Although MgO exhibits a high capacity (exceeding 70 wt %), its regenerability in multicycle tests shows a significant loss of capacity with cycle number prior to lining out. On the other hand, the MgO−Na2CO3 double salt shows a lower (∼16 wt %) but stable capacity over multiple cycles under pressure swing operation. The purpose of this paper is to elaborate on the concept of molten salts as catalysts for CO2 absorption by MgO, and extend these observations to the MgO-containing double salt oxides. We will show that the phenomena involved with CO2 absorption by MgO and MgO-based double salts are similar and general, but with some important differences. This paper focuses on the following key concepts: (i) identification of conditions that favor or disfavor participation of isolated MgO during double salt absorption, and investigation of methods to increase the absorption capacity of double salt systems by including MgO participation; (ii) examination of the relationship between CO2 uptake and melting point of the promoter salt, leading to the recognition of the role of premelting (surface melting) in these systems; and (iii) extension of the reaction pathway model developed for the MgO−NaNO3 system to the double salt systems. This information advances our understanding of MgO-based CO2 absorption systems for application with precombustion gas streams.

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INTRODUCTION The objective of this work is to develop inexpensive and regenerable high capacity warm temperature absorbents (300− 470 °C) for regenerable CO2 capture and release from precombustion syngas streams. An approach employing reversible formation and decomposition of metal carbonates, most notably the MgO and MgCO3 couple, is a leading candidate from a thermodynamic perspective in this temperature range. However, development of a practical, regenerable MgO-based absorbent has proven difficult.1 It appears that kinetics, rather than thermodynamics, is the limiting factor in conversion of MgO.2 Kinetic limitations have recently been described for the one step regeneration of MgCO3 with H2O: MgCO3(s) + H2O → Mg(OH)2(s) + CO2, making a system based on Mg(OH)2 impractical for industrial use.3 In a previous publication we have described the important role of molten salts, notably NaNO3, in facilitating CO2 capture, through a dissolution-recrystallization mechanism, in which kinetic limitations have been removed for the CO2 capture step.4 NaNO3, which plays a catalytic role, can also facilitate the decomposition of MgCO3 back to MgO.5 Molten salts facilitate the capture of CO2 not only by MgO but also by double salts, such as MgO−Na2CO3. Work by Air Products describes several double salt systems comprising MgO combined with Na2CO3, K2CO3, or Li2CO3 in various ratios, prepared by precipitation from magnesium nitrate and soluble carbonate salts, for the capture of CO2.6 The amount of CO2 reversibly captured appears to depend on composition and preparation method. Others have reported difficulty in © 2015 American Chemical Society

reproducing the performance of these materials using the methods described in the Air Products patent, at best having achieved ∼10 wt % reversible capacity with MgO−K2CO3.7 Further examination of this work reveals that residual alkali nitrate salts can remain with the double salt following filtration, and CO2 capture performance can depend on the extent of precipitate washing and thus the residual retained alkali nitrate salt.8 An improved synthesis method, developed in our laboratory, has obviated this reproducibility problem, leading to a more reliable and quantitative examination of these double salt systems. We have described a phase transfer catalysis mechanism to explain the enhanced CO2 absorption on MgO in the presence of molten NaNO3,4 and we submit that the same phenomena can be described in molten salt-promoted double salt systems. In this paper, we examine in further detail these double salts as CO2 absorbers, with an aim to provide additional supporting results and to address several effects that have not been fully characterized and rationalized with these materials. First, we summarize our test results with double salts promoted by several nitrate salts, highlighting a promising and inexpensive MgO−CaCO3 double salt material prepared from natural dolomite; second, we review the difference in double salt capacities relative to MgO itself, the role of temperature in explaining these differences, and conditions under which both Received: November 2, 2015 Revised: December 19, 2015 Published: December 21, 2015 1089

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Figure 1. Cyclic pressure swing test of commercial dolomite with and without 20 wt % NaNO3 at 400 °C.

constant partial pressure of CO2. In the subsequent cycles, the temperature cycled from absorption temperature to desorption temperature. In TPCSA tests, the absorbent was preheated to the desired absorption temperature in CO2 at a heating rate of 7.5 °C/min, and in the subsequent desorption cycle the temperature was raised and the surrounding gas was switched to N2 at the same time to facilitate desorption. The absorption and desorption duration time for each test was determined based on what was needed to obtain stable absorption and regeneration results. More details about the absorption tests are provided in a previous publication.7 Characterization. The phase components of the absorbents were identified by X-ray diffraction (XRD, Bruker D8 ADVANCE) using both standard and in situ measurements, with a scanning rate of 2°/min, using Cu Kα radiation. In one test involving the NaNO3-promoted MgO−Na2CO3 double salt absorbent, in situ XRD measurements were conducted through TSA in a 100% CO2 environment. About 0.5 g of absorbent was loaded for the XRD measurements. A scan was taken at room temperature. Then the absorbent was heated to 380 °C in 100% CO2 and a second scan was taken after holding for 30 min. After the scan was complete, the temperature was increased to 470 °C, and the scan for desorption was taken after holding for 20 min. The same procedures continued for the subsequent absorption and desorption cycles. Similar in situ XRD measurements were also conducted in our previous work, with the main difference being that the absorbent was exposed to CO2 during the ramping up step in this work, whereas N2 was employed previously. Scanning electron microscopy (SEM) analysis was conducted using a FEI Quanta microscope.

MgCO3 and MgNa2(CO3)2 or MgCa(CO3)2 can be formed simultaneously, leading to greater CO2 absorption capacity; third, we compare performance vs melting point of the molten salt, leading to an understanding of the role of premelting (surface melting); and fourth, we discuss and extend the model developed to describe CO2 capture with MgO (plus NaNO3) to the double salt systems.

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EXPERIMENTAL METHODS Material Preparation. Commercially available Mg5(CO3)2(OH)2·4H2O (99%, Sigma-Aldrich) was used in preparing MgO by calcination at 400 °C for 3 h in static air. The obtained MgO, the selected carbonate compound that participates in the double salt formation, and the nitrate salt, in amounts to produce the desired ratios, were mixed together using a ball milling method. In a typical procedure, 10 g of solid absorber components were added to 100 mL of 2-propanol (EMD Chemicals, Canada) along with 192 g of zirconia beads (96 g of beads with a diameter of 1 cm and 96 g of beads with a diameter of 0.3 cm) and milled for 24 h at 60 rpm. The obtained slurry was dried at 60 °C for 12 h to allow the evaporation of 2-propanol. Following drying, the cake was calcined at 500 °C in air for 3 h. Natural mineral dolomite (City Chemical, USA) was also used to prepare the MgO−CaCO3 absorbent. NaNO3 was mixed with dolomite powder employing the ball milling method described above. The obtained slurry was dried, followed by calcination at 450 °C for 3 h. Absorption Tests. The multicycle absorption capacities of the synthesized absorbents were measured using a thermogravimetric analyzer (TGA, Netzsch Thermiche Analyses, STA 409 cell) through pressure swing absorption (PSA), temperature swing absorption (TSA), and temperature−pressure combined swing absorption (TPCSA) at ambient pressure. The weight of the absorbent sample for each test was approximately 20 mg. In the PSA test, the absorbent was preheated to the desired temperature in CO2 at a heating rate of 7.5 °C/min, following which the surrounding gas was switched to N2 for desorption while the temperature was held constant. In the subsequent isothermal cycles, the gas switched between N2 and CO2 for the duration time for each step. In the TSA test, the absorbent was preheated to the desired absorption temperature in CO2 at a heating rate of 7.5 °C/ min, then the temperature was raised for desorption under a

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RESULTS AND DISCUSSION Double Salt Performance and Participation of MgO in Double Salt Absorption. Figure 1 shows the CO2 capture performance of both a MgO−CaCO3 double salt prepared from natural mineral dolomite with 15 wt % NaNO3, and the same double salt in the absence of NaNO3. MgO−CaCO3 is prepared by calcining natural mineral dolomite, CaMg(CO3)2, at 450 °C to convert it into MgO−CaCO3. CaCO3 does not decompose at this temperature. MgO−CaCO3 with NaNO3 shows a regenerative absorption capacity of 15−20 wt %, while MgO−CaCO3 without NaNO3 shows no significant absorption. The dolomite system appears unique in having no 1090

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The double salt systems (Na2Mg(CO3)2 and CaMg(CO3)2) present stable, moderate CO2 absorption capacity, while MgO exhibits a much higher capacity (exceeding 70 wt %) but with some capacity loss over several cycles until lining out of capacity at ∼25 wt %.4 A different CO2 absorption protocol, preheating in CO2 rather than N2, was adopted to investigate the participation of isolated MgO during double salt absorption. Figure 3a shows the CO2 capture performance of a MgO−Na2CO3 double salt prepared with equivalent weights of MgO and Na2CO3 (44 wt %) and 12 wt % NaNO3, providing a sample that is rich in MgO (2.5:1 molar ratio). The figure shows an isothermal pressure swing absorption (PSA) test; i.e., CO2 was flowed during absorption and N2 during desorption, while the temperature was maintained at 400 °C. The duration time for both absorption and desorption was 60 min. The second through ninth cycles show a constant CO2 uptake of around 17 wt %, which is assigned to double salt absorption. The first cycle differs from the rest, showing a capacity of 32.6 wt %. This can be assigned to absorption by both MgO (15.6 wt % CO2 uptake) and MgO−Na2CO3 (17 wt % CO2 uptake). More detailed examination of the figure shows that the majority of the CO2 uptake during the first cycle occurs during the temperature ramp, i.e., where MgO can participate in the absorption. MgCO3 is unstable relative to MgO + CO2 at 400 °C and ambient pressure, thus subsequent absorption cycles carried out isothermally at 400 °C show no effect of the presence of MgO, and only double salt absorption is observed. The MgCO3−Na2CO3 double salt is more stable thermodynamically than MgCO3 by approximately 35 kJ/mol, and as a result its capacity as the fully loaded carbonate is maintained at 400 °C under a CO2 atmosphere.8 These assignments are supported by the X-ray spectra taken at various stages of absorption and desorption of CO2 with this material, as seen in Figure 3b. The phases present with the asprepared absorbent at room temperature include MgO, Na2CO3, and NaNO3. After heating to the desired absorption temperature of 380 °C in CO2, XRD patterns indicate the phases present include Na2Mg(CO3)2, MgCO3, and a small amount of MgO. The NaNO3 phase is not observed, since its melting point (308 °C) is below the measurement temperature, so that NaNO3 melts and does not provide a crystalline diffraction pattern. After desorption at 470 °C, the phase components convert back to Na2CO3 and MgO. After the second absorption, the phase components are identified as Na2Mg(CO3)2 and MgO, while no MgCO3 is observed. This is consistent with the absorption temperature of 400 °C, wherein MgCO3 is unstable whereas the double salt Na2Mg(CO3)2 is stable. We conclude that the observed high initial absorption capacity is due to formation of MgCO3 along with Na2Mg(CO3)2 during the initial temperature ramp. For the MgO−Na2CO3 double salt, a material with a molar ratio of Na2CO3:MgO = 1:1 (mass ratio: 2.5:1) would be expected to give the highest possible double salt absorption capacity, 26.5 wt %, assuming 12 wt % NaNO3. However, this is not the case, and some of the available MgO appears to remain unreacted. This may be the result of the limits in solubility of MgO and/or Na2CO3 in the molten salt. Higher Na2CO3 content leads to rapid degradation over several cycles. Although excess MgO did not participate in the double salt CO2 absorption cycles shown in Figure 3a, it might act as a structure stabilizer. Increasing Absorption Capacity of Double Salt Compositions by Including MgO Participation. Isother-

observable capacity in the absence of the nitrate salt promoter, whereas the MgO−Na2CO3 and MgO−K2CO3 systems show a modest CO2 absorption in the absence of a molten salt promoter, as shown in Table 1. We will discuss further the Table 1. Effect of Different Nitrate Compositions and Concentrations on the Formation of Double Salt Absorbents, with Measurement by PSA at 400 °C carbonate additive

nitrate or carbonate salt

quantity nitrate or carbonate (wt %)

CO2 capacity, 8th cycle (wt %)

sorption product

Metal Oxide (MgO) + Group-I Carbonate + Group-I Nitrate Na2CO3 − 0 3.5 Na2Mg(CO3)2 Na2CO3 NaNO3 2 4.1 Na2Mg(CO3)2 Na2CO3 NaNO3 12 17.2 Na2Mg(CO3)2 Na2CO3 NaNO3 40 0.2 Na2Mg(CO3)2 Na2CO3 LiNO3 12 17.7 Na2Mg(CO3)2 Na2CO3 KNO3 12 17.1 Na2Mg(CO3)2 K2CO3 NaNO3 12 8.4 K2Mg(CO3)2 K2CO3 − 0 3.9 K2Mg(CO3)2 Metal Oxide (MgO) + Group-II Carbonate + Group-I Nitrate CaCO3 NaNO3 15 19.4 CaMg(CO3)2 CaCO3 − 0 0 −

performance of these systems, in terms of premelting (surface melting) behavior, in a later section. Interestingly, MgO− CaCO3 with NaNO3 shows an increasing capacity with cycle number, indicating a self-activating process. The XRD result in Figure 2 verifies that the absorption product is the

Figure 2. X-ray diffraction pattern of CaCO3-promoted MgO absorbent after CO2 absorption.

stoichiometric compound dolomite, CaMg(CO3)2. Dolomite is an abundant naturally available mineral and a NaNO3promoted dolomite material system provides a highly promising warm temperature absorbent candidate. Different double salt systems have different thermodynamic properties, and this can provide a means to develop absorbents tailored for processes with specific operating windows. We mention here that a variant on the dolomite system, dolomite plus K2CO3 (without nitrate salt), showed significant absorption by the MgO component at elevated pressures (20 bar), enhanced further when H2O vapor was also present, although recyclability was not demonstrated.9 In the systems described in this paper, by comparison, water vapor was not present in the feed and the systems were tested at 1 bar pressure. 1091

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Figure 3. (a) Cyclic temperature swing absorption test (380/470 °C) of Na2CO3−MgO absorbent (12 wt % NaNO3), ramping up in CO2. (b) In situ X-ray diffraction patterns of Na2CO3-promoted MgO absorbent during TSA (380/470 °C) absorption cycles, ramping up in CO2.

mal operation with CO2/N2 at 400 °C (PSA mode) shows double salt absorption only, with MgO acting as a spectator (and possibly structure stabilizer) during the absorption/ desorption cycling. The lack of MgO participation is due to unfavorable thermodynamics for absorption at 400 °C. In Figure 4, thermodynamic data for MgCO3 and the double salts are shown, as obtained from HSC (V. 6.1, Qutotec) and Factsage 6.0 when available, or by ab initio computational calculation and experimental verification when not available.8,10,11 According to the thermodynamic data, the trend is that the double salt system can take up CO2 over a higher temperature range than can MgO, whereas the reaction of MgO with CO2 becomes unfavorable at temperatures above 400 °C (PCO2 = 1 atm). There is a difference between the data from HSC and Factsage for MgO, and our experimental data are more consistent with the data from Factsage for MgO. We investigated increasing the capacity of the absorbent by lowering the temperature of operation during the absorption step, so that both MgO and double salt absorption could occur simultaneously. The requirements for simultaneous absorption and regeneration of both compounds suggested dual temperature operation (360 °C absorption, 400 °C desorption) as well as changing the flow gas from CO2 to N2 during desorption (resulting in a temperature−pressure combined swing absorption, TPCSA). Figure 5a shows the results using a

Figure 4. CO2 partial pressure in equilibrium with MgCO3, MgCa(CO3)2, and Na2Mg(CO3)2 as a function of temperature from various sources. Key: (a) HSC; (b) Factsage; (c) experimental measurement by Zhang et al.;8 (d) ab initio calculation by Duan et al.10,11

material having a composition rich in MgO: Na2CO3 (11 wt %); MgO (77 wt %); NaNO3 (12 wt %). It can be seen that a higher capacity is in fact provided by this approach. Initial CO2 capacity at 70.8 wt % dropped to 46.6 wt % by cycle 7. A subsequent extended 30-cycle absorption test of this material 1092

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Figure 5. Cyclic CO2 absorption through combined swing of adsorption at 360 °C in CO2 and desorption at 400 °C in N2 with absorbent composition (a) Na2CO3 (11 wt %), MgO (77 wt %), and NaNO3 (12 wt %) and (b) CaCO3 (11 wt %), MgO (77 wt %), and NaNO3 (12 wt %).

coarsening as the absorption capacity experiences a decrease from the first to seventh cycle, shown in Figure 7, parts a and b. As the absorption capacity decreases, some of the MgO does not participate in the reaction with CO2. Instead, it may act as a structure stabilizer. The absorption capacity eventually reaches a constant value without further significant degradation. Relationship between Molten Salt Melting Point and Onset of CO2 Absorption. We have reported previously that the temperature onset for the capture of CO2 correlates with the melting point of the molten salt. This was demonstrated in an experiment in which MgO powder was added directly on top of a layer of NaNO3 powder, followed by heating the mixture in the presence of CO2 at ambient pressure, using a thermogravimetric analyzer (TGA) to measure uptake. In that case, CO2 absorption began when the temperature reached the melting point of NaNO3, 308 °C.4 A somewhat different result was obtained when the sample was prepared to maximize the interfacial contact between the molten salt and the absorbing oxide (MgO or MgO-containing double salt). Parts a and b of Figure 8 show the CO2 uptake versus temperature for the systems NaNO3−MgO−Na2CO3 and NaNO3−MgO, respectively, which had been thoroughly ball milled and calcined at 400 °C to allow melting of NaNO3. Significant uptake of CO2 is seen to begin at 250 °C with NaNO3−MgO, and with the double salt a small amount of uptake can even be seen initiating near 210 °C, clearly well below the melting point of the additive salt. This behavior is consistent with the phenomenon “premelting” or “surface melting”. Frenken and van der Veen

showed that the absorption capacity eventually lined out slightly above 25 wt %, as seen in Figure 6. This capacity is

Figure 6. 30-cycle CO2 absorption through combined swing of adsorption at 360 °C in CO2 and desorption at 400 °C in N2 with absorbent composition Na2CO3 (11 wt %), MgO (77 wt %), and NaNO3 (12 wt %).

higher than the double salt alone. A similar approach was utilized to increase absorption capacity for the MgO−CaCO3 system with a composition of CaCO3 (11 wt %), MgO (77 wt %), and NaNO3 (12 wt %), as shown in Figure 5b. The absorption experiences an initial self-activation process (capacity increase), but is similarly followed by degradation from 62.7 wt % to 46.1 wt % by cycle 7. Morphology examination of the MgO−Na2CO3 material indicates particle

Figure 7. Morphology of absorbent with composition Na2CO3 (11 wt %), MgO (77 wt %), and NaNO3 (12 wt %) after (a) 1st absorption and (b) 7th absorption. 1093

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Figure 8. CO2 absorption as a function of temperature for a sample comprising (a) 77 wt % MgO, 11 wt % Na2CO3, and 12 wt % NaNO3 and (b) 80 wt % MgO and 20 wt % NaNO3. Both show CO2 absorption at a temperature below the melting point of NaNO3 (308 °C).

define “premelting” as a “gradual disordering of the surface region”, and “surface melting” as “the presence of a surface liquid film on top of, and in equilibrium with, a well ordered substrate”.12 Premelting therefore precedes surface melting. Yang et al. define premelting as “the formation of thermodynamically stable liquid films at solid interfaces subjected to temperatures below but near the bulk melting temperature”.13 We choose to employ the latter definition. Regardless of terminology, the effect is driven by thermodynamics: a thin film of metastable liquid forms as a result of lowering of total interfacial free energy relative to the solid− solid interface. This effect was investigated further. A bulk piece prepared from a ball milled sample comprising MgO and NaNO3 was put in the TGA/DSC unit and then covered with extra NaNO3 solid, such that when the latter melted it would fully cover the surface of the MgO, as shown in Figure 9a. Figure 9b provides the DSC and TGA traces when this sample was heated in the presence of CO2. The DSC trace shows evidence of a premelting event, which correlates with the initiation of weight

gain (see the corresponding TGA trace). The DSC trace also shows a larger peak corresponding to the melting of NaNO3 at 308 °C, however the TGA trace shows no weight gain when the temperature reaches or rises above the melting point. This can be explained by the complete coverage of MgO by NaNO3 once it has melted, and no CO2 can be absorbed because CO2 transport through the molten NaNO3 to the MgO surface is not a viable pathway. This result is consistent with the assertion made previously that CO2 initially adsorbs onto bare MgO and then migrates to the triple phase boundary.4 Figure 9b demonstrates that premelting of the nitrate salt clearly occurs, forming a thin surface film of NaNO3 prior to reaching its bulk melting point.12−16 Although not as efficient as a fully melted salt in facilitating CO2 capture, premelted NaNO3 appears to be sufficient to promote CO2 absorption at 250 °C by MgO. We speculate that the further decrease in the onset absorption temperature when Na2CO3 is also present is the result of additional surface interactions favoring reduction in the temperature of onset for premelting. In order to explain the absorption of CO2 on NaNO3-free systems such as MgO−Na2CO3, premelting must nevertheless be occurring at the solid−solid interface, facilitating formation of accessible Mg2+ cations that can react with CO2. Differences in interfacial energies may explain the differences in the absorption behavior among MgO−Na2CO3, MgO−CaCO3, and MgO−K2CO3 summarized in Table 1. The dolomite system appears unique in having no observable capacity in the absence of the nitrate salt promoter, whereas the MgO− Na2CO3 and MgO−K2CO3 systems show some modest CO2 absorption. Similarly, CO2 absorption by the Cs2CO3−MgO double salt occurs without other low melting salt additives.17 It is worth noting that the melting points of the carbonate components contained in the four absorbent systems rank in the order of CaCO3 > K2CO3 > Na2CO3 > Cs2CO3. All melt above the operating temperature of the experiment, so that premelting must be occurring. CaCO3 has the highest melting point and the largest difference between melting point and the operating temperature of the absorption measurement. This may explain the lack of a premelting effect in the dolomite system in absence of the molten nitrate salt additive. Mechanism for CO2 Absorption by Double Salts. We have shown that the capture of CO2 by MgO and MgO-based double salts, facilitated by salt additives that become molten at the operating temperature, is a general phenomenon. For MgO

Figure 9. (a) NaNO3 solids are added on top of a MgO + NaNO3 bulk sample, sufficient to fully cover the sample when melted. (b) TGA and DSC curves for the system depicted in Figure 5a, showing result of ramping the temperature during exposure to CO2. 1094

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Na2CO3 or MgCO3−CaCO3, provided the temperature is sufficiently high that MgCO3 is thermodynamically unstable. At 400 °C under PSA operation, a stable capacity of the double salt of ∼16 wt % is obtained. Lower temperature absorption (below 375 °C) allows both free MgO and the MgO that is complexed as the double salt (MgO−Na2CO3 or Mg−CaCO3) to capture CO2 simultaneously. In this latter case, a higher overall capacity can be achieved than by the double salt alone. Operating to gain this higher capacity is best obtained through both a temperature increase and gas switching (to N2) during regeneration. This potential for high capacity provides an incentive for further study and optimization. MgO−CaCO3 can be prepared as an effective sorbent from the natural mineral dolomite, an abundant natural compound, thereby providing an additional incentive for implementation through low cost raw materials. Detailed examination of the absorption/desorption behavior shows that CO2 absorption can occur at a temperature below the true melting point of the salt. This premelting behavior has been described in the literature, but this is the first report of its occurrence in a CO2 capture system. For MgO alone, premelting is thought to occur through forming a thin film of the salt onto the MgO surface, below its melting point, which facilitates some dissolution of MgO. The magnitude of this effect, i.e. how much the capture temperature is lowered, is a function of the interfacial free energies between MgO, premelted salt, and bulk salt relative to the interfacial free energy between MgO and the bulk salt. The mode of operation for CO2 capture that has been described previously with MgO and NaNO3 can be applied by extension to the double salt systems. The important criterion is the ability of the molten salt to partially dissolve the MgO and the alkali carbonate components, allowing them to react with CO2. Since high concentrations of molten salt decrease the observed capacity, direct CO2 absorption into the molten salt phase can be discounted. More consistent with the data is an initial physisorption by CO2 onto the uncovered solid MgO surface, followed by its migration to the triple phase boundary (solid, melt, gas phase CO2). There it reacts with the partially dissolved MgO present in the melt and precipitates as the double salt product. It is also possible to absorb CO2 in the alkali metal double salt systems (MgO plus Na2CO3, K2CO3, or Cs2CO3, but not CaCO3) without the presence of a melting salt such as NaNO3. Formation of a premelt must also be occurring in these systems, but capacity is lower than when the NaNO3 is present, most likely the result of a less effective premelt phase.

as a single component absorbent, the promoting effect is consistent with the partial dissolution of solid MgO into the molten salt. We assert that the molten salt acts as a type of phase transfer catalyst.4 Through this step, the high lattice energy constraints, which limit direct reaction between MgO and CO2, are removed. At high loadings of the nitrate salt, CO2 absorption is again hindered (Table 1), which indicates that a mechanism by which CO2 dissolves into the molten salt, and is transported through the melt to the MgO surface for reaction, can be discounted. A certain amount of available MgO surface must be required for surface adsorption of CO2, and this physisorbed CO2 migrates to the triple phase boundary, where it reacts with dissolved Mg2+ to form MgCO3. The point at which MgCO3 begins to precipitate is determined by its solubility in the molten salt. Different solubilities can explain the differences in performance of the different salts as molten phase catalysts. A similar pathway can explain the promoting effect in the double salt systems, as shown in Figure 10. Double salt

Figure 10. Conceptual diagram showing CO2 absorption and reaction with MgO in the presence of Na2CO3 to form MgNa2(CO3)2 double salt at the triple phase boundary.

absorbent materials contain MgO, the carbonate component (such as Na2CO3 or CaCO3), and the salt additive that becomes a molten phase during absorbent operation. Similar to the metal oxide absorbent system, MgO in the double salt absorbent dissolves in the molten salt and provides solvated [Mg2+···O2−] ionic pairs, which have much weaker interaction than the strong Mg−O ionic bonds in bulk MgO. The carbonate component in the double salt absorbent, such as Na2CO3 as shown in Figure 10, can partially dissolve in the molten salt in similar manner into ionic pairs. Carbonate salts are found to have appreciable solubility in molten salts.18,19 With the dissolution of solid MgO and the carbonate salt, the solid reactants are transferred to a liquid medium where they gain enhanced mobility. They readily combine with CO2 at the triple phase boundary to form the double salt carbonate, which precipitates from the liquid medium as it reaches its solubility limit. By operating at a temperature (400 °C) where MgCO3 will not precipitate, solid formation only occurs through double salt precipitation. Premelting of the nitrate salt occurs in these systems, so that CO2 absorption can occur at temperatures below the melting point of the salt additive. However, absorption during premelting generally is less efficient than under conditions where the fully molten salt is present.

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AUTHOR INFORMATION

Corresponding Author

*(D.L.K.) E-mail: [email protected]. Telephone: (509) 375-3908. Fax: (509) 375-2186. Present Addresses §

(K.Z.) 6752 Baymeadow Drive, Glen Burnie, MD 21060. Email: [email protected]. Telephone: (401) 327-1014. ⊥ (H.C.)School for Engineering of Matter, Transport and Energy, Arizona State University, Tempe, AZ 85281. E-mail: [email protected].

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CONCLUSIONS Salt additives that are present as a molten phase are effective in facilitating CO2 capture by alkaline earth oxides, primarily MgO. With MgO, admixing with Na2CO3 or CaCO3 allows the sole formation of a double salt upon capture of CO2: MgCO3−

Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. K.Z. conducted the majority of experimental study and data analysis. X.S.L. conducted some of the 1095

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(13) Yang, Y.; Asta, M.; Laird, B. B. Solid-Liquid Interfacial Premelting. Phys. Rev. Lett. 2013, 110, 96−102. (14) Janz, G. J.; Kelly, J.; Perano, J. L. Melting and Pre-Melting Phenomena in Alkali Metal Nitrates. J. Chem. Eng. Data 1964, 9, 133− 136. (15) Clarke, D. R. On the Equilibrum Thickness of Integranular Glass Phases in Ceramic Materials. J. Am. Ceram. Soc. 1987, 70, 15−22. (16) Luo, J.; Chiang, Y.-M. Wetting and Prewetting on Ceramic Surfaces. Annu. Rev. Mater. Res. 2008, 38, 227−249. (17) Liu, M.; Vogt, C.; Chaffee, A. L.; Chang, S. L. Y. Nanoscale Structural Investigation of Cs2CO3-Doped MgO Sorbent for CO2 Capture at Moderate Temperature. J. Phys. Chem. C 2013, 117, 17514−17520. (18) FACT-Web Programs. http://www.crct.polymtl.ca/factweb. php. (19) Cherginets, V. L.; Deineka, T. G.; Demirskaya, O. V.; Rebrova, T. P. Potentiometric Investigation of Oxide Solubilities in Molten KClNaCl Eutectic. J. Electroanal. Chem. 2002, 531, 171−178.

experimental study and participated in all data analysis. D.L.K. provided project leadership and manuscript review. H.C. conducted some of the preliminary experimental studies. P.S. provided technical input. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support from the US DOE Office of Fossil Energy (NETL), the State of Wyoming, and PNNL internal investment (LDRD-ECI) is gratefully acknowledged. Some work was carried out at PNNL’s Environmental and Molecular Science Laboratory (EMSL), a national scientific user facility sponsored by the DOE’s office of Biological and Environmental Research (BER). Helpful discussions with Wei-Zhen Li are gratefully acknowledged.

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ABBREVIATIONS TGA, thermogravimetric analyzer; PSA, pressure swing absorption; TSA, temperature swing absorption; TPCSA, temperature−pressure combined swing absorption; XRD, Xray diffraction; SEM, scanning electron microscopy; TPB, triple phase boundary; EDS, energy dispersive X-ray spectroscopy

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REFERENCES

(1) Choi, S.; Drese, J. H.; Jones, C. W. Absorbent Materials for Carbon Dioxide Capture from Large Anthropogenic Point Sources. ChemSusChem 2009, 2, 796−854. (2) Gregg, S. J.; Ramsay, J. D. Absorption of Carbon Dioxide by Magnesia Studied by Use of Infrared and Isotherm Measurements. J. Chem. Soc. A 1970, 2784. (3) Fisher, J. C.; Siriwardane, R. V. Mg(OH)2 for CO2 Capture from High-Pressure, Moderate-Temperature Gas Streams. Energy Fuels 2014, 28, 5936−5941. (4) Zhang, K.; Li, X. S.; Li, W.-Z.; Rohatgi, A.; Duan, Y.; Singh, P.; Li, L.; King, D. L. Phase Transfer-Catalyzed Fast CO2 Absorption by MgO-Based Absorbents with High Cycling Capacity. Adv. Mater. Interfaces 2014, 1, 1400030. (5) Zhang, K.; Li, X. S.; Li, W.-Z.; Rohatgi, A.; Duan, Y.; Singh, P.; Li, L.; King, D. L. Phase Transfer-Catalyzed Fast CO2 Absorption By MgO-Based Absorbents With High Cycling Capacity. Adv. Mater. Interfaces 2014, 1, 1400030 see supplementary material.. (6) Mayorga, S. G.; Weigel, S. J.; Gaffney, T. R., Brzozowski, J. R. Carbon Dioxide Adsorbents Containing Magnesium Oxide Suitable for Use at High Temperatures. 2001, US 6,280,503 Bl. (7) Xiao, G.; Singh, R.; Chaffee, A.; Webley, P. Advanced Adsorbents Based on MgO and K2CO3 for Capture of CO2 at Elevated Temperatures. Int. J. Greenhouse Gas Control 2011, 5, 634−639. (8) Zhang, K.; Li, X. S.; Duan, Y.; King, D. L.; Singh, P.; Li, L. Roles of Double Salt Formation and NaNO3 in Na2CO3-Promoted MgO Absorbent for Intermediate Temperature CO2 Removal. Int. J. Greenhouse Gas Control 2013, 12, 351−358. (9) Zarghami, S.; Hassanzadeh, A.; Arastoopour, H.; Abbasian, J. Effect of Steam on the Reactivity of MgO-Based Sorbents in Precombustion CO2 Capture Processes. Ind. Eng. Chem. Res. 2015, 54, 8860−8866. (10) Duan, Y.; Sorescu, D. C. CO2 Capture Properties of Alkaline Earth Metal Oxides and Hydroxides: A Combined Density Functional Theory and Lattice Phonon Dynamics Study. J. Chem. Phys. 2010, 133, 074508. (11) Duan, Y.; Zhang, B.; Sorescu, D. C.; Johnson, J. K. CO2 Capture Properties of M−C−O−H (M = Li, Na, K) Systems: A Combined Density Functional Theory and Lattice Phonon Dynamics Study. J. Solid State Chem. 2011, 184, 304−311. (12) Frenken, J. W. M.; Veen, J. F. v. d. Observation of Surface Melting. Phys. Rev. Lett. 1985, 54, 134−137. 1096

DOI: 10.1021/acs.jpcc.5b10729 J. Phys. Chem. C 2016, 120, 1089−1096