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International Journal of Greenhouse Gas Control 12 (2013) 351–358

Contents lists available at SciVerse ScienceDirect

International Journal of Greenhouse Gas Control

j ourna l ho mepage: www.elsev ier .com/ locate / i jggc

oles of double salt formation and NaNO3 in Na2CO3-promoted MgO absorbentor intermediate temperature CO2 removal

eling Zhanga,b, Xiaohong S. Li c, Yuhua Duand, David L. Kingc,∗, Prabhakar Singha,b, Liyu Li c

Center for Clean Energy Engineering, University of Connecticut, Storrs, CT 06269, USADepartment of Chemical, Materials and Biomolecular Engineering, University of Connecticut, Storrs, CT 06269, USAInstitute for Integrated Catalysis, Pacific Northwest National Laboratory, P. O. Box 999, Richland, WA 99354, USANational Energy Technology Laboratory, United States Department of Energy, Pittsburgh, PA 15236, USA

r t i c l e i n f o

rticle history:eceived 14 May 2012eceived in revised form 5 October 2012ccepted 12 November 2012

eywords:O2 absorption and desorptiona2CO3 promoted MgOa2Mg(CO3)2 double saltarm temperature CO2 capture

a b s t r a c t

Absorption and desorption of carbon dioxide on Na2CO3-promoted MgO have been studied at temper-atures compatible with warm gas cleanup (300–470 ◦C) from a pre-combustion syngas. The absorbentsare synthesized through the formation and activation of the precipitate resulting from the addition ofsodium carbonate to an aqueous solution of magnesium nitrate. The absorbent, which comprises MgO,Na2CO3 and residual NaNO3 after activation, forms the double salt Na2Mg(CO3)2 on exposure to CO2.The thermodynamic properties of the double salt, obtained through computational calculation, predictthat the preferred temperature range for absorption of CO2 with the double salt is significantly highercompared with MgO. Faster CO2 uptake can be achieved as a result of this higher temperature absorp-tion window. Absorption tests indicate that the double salt absorbent as prepared has a capacity toward

re-combustion CO2 captureaNO3 facilitated CO2 capture

CO2 of 15 wt.% (3.4 mmol CO2/g absorbent) and can be easily regenerated through both pressure swingand temperature swing absorption in multiple-cycle tests. Thermodynamic calculations also predict animportant effect of CO2 partial pressure on the absorption capacity in the warm temperature range. Theimpurity phase, NaNO3, is identified as a key component in facilitating CO2 absorption by these materials.The reason for reported difficulties in reproducing the performance of these materials can be traced tospecific details of the synthesis method, which are reviewed in some detail.

. Introduction

Fossil-fueled power plants are by far the largest CO2 emittersnd major contributors to greenhouse gas emissions, making thembvious targets for the implementation of advanced carbon dioxideapture and storage technologies. Three predominant technologieshat produce CO2 from coal power plants are oxy-combustion, post-ombustion and pre-combustion (Herzog, 2009). Pre-combustionroduces syngas, which may be used in a number of processes,uch as methane, methanol, and Fischer–Tropsch synthesis, and2 production for fuel cell power production or for IGCC applica-

ions. In general, syngas must be cleaned of impurities such as sulfurases prior to utilization, and their removal at warm tempera-ures (300–500 ◦C) provides an efficiency improvement in avoidingooling and reheating the syngas prior to use. However, if CO2 is

lso to be captured for subsequent sequestration, warm CO2 cap-ure prior to syngas utilization is also necessary to maintain thefficiency advantage. In some applications, an added benefit is to

∗ Corresponding author. Tel.: +1 509 375 3908; fax: +1 509 375 2186.E-mail address: David.King@pnnl.gov (D.L. King).

750-5836/$ – see front matter © 2012 Elsevier Ltd. All rights reserved.ttp://dx.doi.org/10.1016/j.ijggc.2012.11.013

© 2012 Elsevier Ltd. All rights reserved.

employ the CO2 capture step with a catalytic reaction such as watergas shift, to increase equilibrium conversion. This is only possi-ble with warm CO2 capture in which the two functions operatein the same temperature range (Hufton et al., 2000; Sircar et al.,1995). Most commercial processes for capturing CO2 use alkalinesolutions such as various alkylamines, or physical solvents such asglycol ethers (Blamey et al., 2010; Steeneveldt et al., 2006). Unlikeliquid absorbents, solid sorbents can in principle be used over awider temperature range, from ambient temperature to 700 ◦C.Unmodified MgO has a very low capacity of 0.24 mmol/g at 200 ◦C(a preferred temperature for CO2 absorption based on thermody-namic considerations) (Gregg and Ramsay, 1970), indicating poorabsorption kinetics at that temperature. However, higher tempera-ture operation is limited by thermodynamic equilibrium: accordingto HSC Chemistry (V. 6.1, Qutotec), MgCO3 decomposes to MgOand CO2 above 300 ◦C (at 1 bar CO2 pressure). At higher pressures,higher operating temperatures are possible and the performance ofMgO could improve, but the observed low kinetic rates and capaci-

ties remain a concern. Recently, several MgO-based materials withsignificantly better performance have been reported to selectivelyand reversibly absorb CO2. An absorption capacity of 3.37 mmol/gwas reported for Mg(OH)2, however, the operation of this sorbent is

3 f Greenhouse Gas Control 12 (2013) 351–358

lSasrto3mfTo4esr8Cam

cbtooats

2

2

wtdplNatTtobrtcff

2

azsatTvns

Fig. 1. X-ray diffraction patterns of Na2CO3-promoted MgO absorbent: (a) afterfiltration; (b) after drying; (c) after activation.

52 K. Zhang et al. / International Journal o

imited to the temperature range of 200–315 ◦C (Siriwardane andtevens, 2009) and requires rehydroxylation of MgO to regener-te the sorbent. A U.S. patent describing MgO-based double saltorbents, also described as alkali promoted MgO-based sorbents,eports a broad capacity range of 1.1–12.9 mmol/g depending onhe conditions of synthesis, with a highest regenerable capacityf 11 mmol/g demonstrated using pressure swing regeneration at75 ◦C (Mayorga et al., 2001). Double salts are salts containingore than one cation or anion, obtained by combining two dif-

erent salts which are crystallized in the same regular ionic lattice.here are stochiometric and non-stoichiometric double salts. In thepen literature, a Na–Mg double salt absorbent with a capacity of.7 mmol/g at 375 ◦C has been described (Singh et al., 2009). How-ver, the authors noted having difficulty in producing reproducibleamples. A recent paper by Xiao et al. describes a double salt mate-ial based on MgO plus K2CO3 that had a maximum capacity of.69 wt.% (∼2 mmol/g) as measured by TGA at 375 ◦C using a pureO2 stream at atmospheric pressure (Xiao et al., 2011). A regener-ble capacity (N2 as the purge gas) of 7.69 wt.% was reported. Thisaterial showed a substantially lower capacity at 400 ◦C.Given the attractive operating temperature range and potential

apacity of the double salt materials for CO2 capture and release,ut with differing reports regarding performance, we determinedo investigate the MgO-based material further. Here we report onur investigations aimed at developing an increased understandingf the key phases present and their transformations upon capturend release of CO2. This in turn will provide knowledge leadingo identifying directions for improving the reproducibility of theample synthesis and optimizing CO2 capture performance.

. Experimental

.1. Materials synthesis

Na2CO3-promoted MgO sorbents were synthesized through aet-chemistry route described by Mayorga et al. (2001). At room

emperature, 3.01 g of Na2CO3 (99.95%, Sigma–Aldrich, USA) pow-er was gradually added over 5 min to a rapidly stirred solutionrepared by dissolving 2.43 g of Mg(NO3)2·6H2O (99.0%, Fluka Ana-

ytical, Germany) in 30 ml deionized water. The reactant ratio ofa2CO3 to Mg(NO3)2·6H2O is approximately 3:1, corresponding tobout 50% molar excess of Na2CO3 relative to that required to formhe stoichiometric double salt. A white slurry formed immediately.he stirring continued for 1 h, and then the mixture was allowedo settle for 24 h. The precipitate was separated by filtration with-ut additional water washing. Both the Na2CO3 in excess and theyproduct NaNO3 have high solubility in water (215 g/L and 912 g/L,espectively). Certain amounts of these two components remain inhe wet cake after filtration, mainly in the retained water. The wetake was oven dried at 120 ◦C for 16 h, and then activated at 400 ◦Cor 3 h in air, with a 5 ◦C/min heating and cooling rate. The yieldrom this synthesis was 0.7–0.8 g.

.2. Absorption tests

The multi-cycle absorption capacity of the synthesizedbsorbent was measured using a thermogravimetric analyzer (Net-sch Thermiche Analyse, STA 409 cell) both through temperaturewing absorption (TSA) and pressure swing absorption (PSA) atmbient pressure. The weight of the absorbent sample for eachest was approximately 20 mg. Operating temperatures for both

SA and PSA were determined by conducting initial trial tests atarious temperatures. The temperatures which provided a combi-ation of high capacity and fast absorption–desorption rates wereelected. Based on the initial test results, an absorption temperature

Fig. 2. SEM image of Na2CO3-promoted MgO absorbent.

K. Zhang et al. / International Journal of Greenhouse Gas Control 12 (2013) 351–358 353

bsorp

oThbtcN3ttChshmov

Fig. 3. First 3 cycles of (a) TSA and (b) PSA CO2 a

f 380 ◦C and a desorption temperature of 470 ◦C were selected forSA tests. Similarly, 400 ◦C was selected for the PSA tests. The initialeating from room temperature to the absorption temperature foroth TSA and PSA tests was conducted in 100% N2 to avoid absorp-ion before reaching the desired temperature. For TSA, during eachooling step from 470 ◦C to 380 ◦C, the surrounding gas was 100%2. The remaining steps during the temperature swing between80 ◦C and 470 ◦C were in 100% CO2. The absorption and desorp-ion time durations were 60 min and 10 min, respectively. The PSAests were carried out by exposing the sample to alternating 100%O2 for 30 min and 100% N2 for 60 min at 400 ◦C. The absorptioneat was measured along with the TG tests through differentialcanning calorimetry (DSC). In order to calibrate the device, the

eat of fusion of high purity NaNO3 (99.995%, Sigma–Aldrich) waseasured by heating to 400 ◦C in N2, and a correction factor was

btained which was applied to the measured CO2 heat of absorptionalue.

tion tests of Na2CO3-promoted MgO absorbent.

2.3. Characterization

The phase components of the absorbents were identified byX-ray diffraction (Bruker D8 ADVANCE) using both standard andin situ measurements, with a scanning rate of 2◦/min, using Cu K�radiation. CO2 absorption during in situ XRD measurements wasconducted through temperature swing between 380 ◦C and 470 ◦Cin a 100% CO2 environment. Scanning electron microscopy (SEM)analysis was conducted using a JEOL JSM-5900LV microscope.

3. Results and discussion

3.1. Absorbent chemistry during synthesis

During the preparation of the absorbent, the phase compo-nents were identified by XRD: after filtration in the form of wetpaste; after drying in the form of a dry cake; and after activation

354 K. Zhang et al. / International Journal of Green

Fp

ipsbOdoidNdNt

N

adWpcmorpt

3

caiTtwarTtha

t4p(pi

ig. 4. 10-Cycle TSA and PSA CO2 absorption capacity comparison for Na2CO3-romoted MgO absorbent.

n the form of a powder. The results are shown in Fig. 1. The wetaste is seen to consist mainly of dypingite (Mg5(CO3)4(OH)2·5H2O,hown in Fig. 1(a). There are unidentified peaks that can likelye attributed to other hydrated forms of the precipitated salts.ther components such as Na2CO3 and NaNO3 are not observedue to their being in a dissolved state. The dried cake is composedf Mg5(CO3)4(OH)2·4H2O, NaNO3, and Na2Mg(CO3)2, as shownn Fig. 1(b). Dypingite loses some of its crystalline water duringrying and becomes Mg5(CO3)4(OH)2·4H2O. As water evaporates,aNO3 precipitates out as crystals and is observed in the X-rayiffraction patterns. Surprisingly, Na2CO3, is not observed; rather,a2Mg(CO3)2 is found, which indicates that the reaction in equa-

ion (a) takes place during drying and consumes the Na2CO3.

a2CO3 + Mg5(CO3)4(OH)2·xH2O → Na2Mg(CO3)2 (a)

Fig. 1(c) shows that the phase components of the absorbentfter activation are MgO, Na2CO3, and NaNO3, which indicates theecomposition of both Mg5(CO3)4(OH)2·4H2O and Na2Mg(CO3)2.ith a melting point of 308 ◦C, NaNO3 experiences a melt-solidify

rocess during activation. Molten NaNO3 has a creep tendency andan spread out instantly and cover the surface of many oxides andetals (Nissen and Meeker, 1983). The smooth morphology of some

f the material observed by SEM (Fig. 2) is the result of melted ande-solidified NaNO3. However, only a partial fraction of the com-onents in the activated absorbent mixture is covered by NaNO3;here are also uncovered coarse surfaces present.

.2. Evaluation of CO2 absorption

The TGA results from the first three absorption–desorptionycles of 10-cycle PSA and TSA tests with Na2CO3-promoted MgObsorbent are shown in Fig. 3. The 10-cycle PSA and TSA capac-ties of Na2CO3-promoted MgO absorbent are presented in Fig. 4.his demonstrates that the Na2CO3 promoted MgO absorbent func-ions in both TSA and PSA modes. For both processes, the absorbenteight increased by 15% over multiple cycles, representing a CO2

bsorption capacity of 3.4 mmol/g. The initial capacity remainedelatively stable, with a slight degradation over the next nine cycles.he absorbent tested through TSA experienced more degradationhan when it was tested through PSA. This is probably due to theigher regeneration temperature for TSA than that for PSA, causing

small amount of structural degradation of the absorbent.The dynamic phase changes occurring during the absorp-

ion cycles are shown in Fig. 5, using TSA between 380 ◦C and70 ◦C. The XRD patterns show that the stoichiometric double salt

hase, Na2Mg(CO3)2, appears during absorption and disappearsdecomposes) during desorption. Correspondingly, the Na2CO3eak strength decreases to a very low level after absorption, show-

ng that most of it is consumed through Na2Mg(CO3)2 formation,

house Gas Control 12 (2013) 351–358

and it increases after regeneration, showing that CO2 is released.We conclude that CO2 absorption proceeds through the reversibleformation of Na–Mg double salt following the reaction (b):

MgO + Na2CO3 + CO2(g) ↔ Na2Mg(CO3)2 (b)

It can be seen from Fig. 5 that there is excess MgO and insuffi-cient Na2CO3 for additional double salt formation. Since the Na2CO3is consumed during CO2 absorption through Na2Mg(CO3)2 forma-tion, its amount can be calculated through the observed absorptioncapacity. NaNO3 salt begins to decompose into NaNO2 at around500 ◦C, and its complete decomposition into Na2O was observedafter 800 min at 649 ◦C (Bauer et al., 2011; Freeman, 1956). In orderto speed up the complete decomposition and measure the NaNO3amount, the absorbent mixture was heated and the decompositioncompleted at 750 ◦C. By measuring the mass loss, the amount ofNaNO3 contained in the absorbent was calculated to be ∼12 wt.%.The mass balance calculation shows that this absorbent contains36 wt.% Na2CO3, 52 wt.% MgO and 12 wt.% NaNO3. The amount ofMgO in the product is consistent with the amount of the startingmaterial Mg(NO3)2·6H2O. It should also be noted that the peak shiftat different temperatures is due to lattice expansion, and it is clearlyobserved in the case of Na2CO3. The NaNO3 phase cannot be seenduring the absorption process since it melts and loses its crystalstructure and becomes undetectable by X-ray diffraction.

Based on the X-ray spectral results, it can be concluded thata fraction of the MgO combines with Na2CO3 to form the dou-ble salt during the absorption process. The remaining MgO, whichdoes not have available Na2CO3 in its proximity, remains as MgO.It displays virtually no absorption capacity under the given con-ditions, i.e., no MgCO3 peak is identified in the absorption cyclesin Fig. 5. As shown in Fig. 1(b), Mg exists in two different forms inthe precursor: Na2Mg(CO3)2 and Mg5(CO3)4(OH)2·4H2O. After acti-vation, both Na2Mg(CO3)2 and Mg5(CO3)4(OH)2·4H2O decompose.Na2Mg(CO3)2 decomposes into Na2CO3 and MgO, however, somevestige of the Na–Mg double salt structure must be maintained,perhaps via the proximity of the components. We suggest that thislocalized structure in the absorbent enables Na2Mg(CO3)2 to easilyform again during CO2 absorption. The conversions of the differ-ent compounds involved in the absorption cycles are illustrated inFig. 6.

The theoretical CO2 absorption capacity for Na2CO3-promotedMgO when producing Na2Mg(CO3)2 is 29.3 wt.% (6.5 mmol/g). Thisis much lower than the reported 11 mmol/g for the best absorbentreported by Mayorga et al. following the same synthesis proce-dures (Mayorga et al., 2001). It is suggested in that work that anon-stoichiometric double salt compound forms during absorp-tion, which would imply that the excess MgO can also react withCO2 and as a result increase the capacity. However, we do notobserve this effect under our test conditions. Although the capac-ity that we report is lower than the reported maximum capacityof 11 mmol/g, this capacity is still high in comparison with otherreported literature data for MgO-based absorbents (Hassanzadehand Abbasian, 2010; Singh et al., 2009; Siriwardane and Stevens,2009; Xiao et al., 2011).

These results indicate that the amount of retained Na2CO3 in theinitial synthesis step directly affects the performance. The amountof Na2CO3 retained in the filter cake is difficult to control, and thisbecomes one of the reasons for the sensitive nature of the synthesistechnique.

3.2.1. Thermodynamic analysisA thermodynamic analysis of the primary phases involved

in this system has been carried out, with thermodynamicvalues provided by HSC Chemistry (V. 6.1, Qutotec) whenavailable, and by ab initio computational calculations whennot available, as is the case for Na2Mg(CO3)2. The details

K. Zhang et al. / International Journal of Greenhouse Gas Control 12 (2013) 351–358 355

ted M

os2

tatwfpNacat

M

MetittT

Fig. 5. In situ X-ray diffraction patterns of Na2CO3-promo

f the computational methods have been described and pre-ented elsewhere (Duan and Sorescu, 2010; Duan et al.,011).

Calculated enthalpy and free energy changes for CO2 cap-ure reactions by MgO, Na2O, and MgO + Na2CO3 (double salt)re shown in Fig. 7(a) and (b), respectively. It can be seen thathe double salt is thermodynamically more stable than MgCO3,ith both a lower enthalpy (by ∼24 kJ/mol) and lower Gibbs

ree energy of formation (by ∼35 kJ/mol) over the warm tem-erature range of interest. As a result, the driving force fora2Mg(CO3)2 formation is larger than MgCO3 formation at equiv-lent conditions. The Gibbs free energy changes of these CO2apture reactions, provided in Fig. 7(b), show the temperaturet which �G = 0 at 1 bar CO2 pressure. MgO can capture CO2o form MgCO3 up to 300 ◦C through reaction (c):

gO + CO2(g) = MgCO3 (c)

Above 300 ◦C, thermodynamically it is more favorable forgCO3 to dissociate to MgO. In the presence of Na2CO3, how-

ver, as described in reaction (b), the thermodynamic data predicthat MgO can capture CO2 at temperatures up to 520 ◦C by form-

ng double salt Na2Mg(CO3)2, due to its greater stability. We notehat in our TSA experiments, we employed 470 ◦C for desorp-ion, which should not be possible as predicted by computation.hus, we see some disparity between theory and experiment. As

Fig. 6. The illustration of the conversions of compounds

gO absorbent during TSA (380/470 ◦C) absorption cycles.

shown in Fig. 7(b) Na2CO3 contained in the synthesized sorbentmixture will remain stable in its carbonate form. Though not con-tributing to CO2 absorption by itself, it activates MgO throughNa2Mg(CO3)2 formation, thereby increasing the effective operatingtemperature range for CO2 capture.

3.2.2. Absorption heat measurementHeat changes associated with the absorption reactions

were measured through differential scanning calorimetric (DSC)analysis in conjunction with the thermogravimetric measure-ments. The DSC peak area associated with each absorption anddesorption process is proportional to the change in enthalpy, theheat consumed or released by the sample. At 400 ◦C, the absorptionheat, which is also the formation heat of Na2Mg(CO3)2, was calcu-lated by taking an average of the DSC results from 10 absorptiontests, yielding a value of −122.4 kJ/mol. We estimate this value to beaccurate within ±5%. According to Fig. 7(a), the theoretical enthalpyfor Na2Mg(CO3)2 formation is −121.6 kJ/mol at 400 ◦C. The exper-imental data is in accordance with theoretical predictions withinthe uncertainty of the measurement.

3.3. Equilibrium pressure study of the absorbent

With the calculated thermodynamic data for Na2Mg(CO3)2,the predicted equilibrium of reaction (b) as a function of both

involved in absorption and desorption processes.

356 K. Zhang et al. / International Journal of Greenhouse Gas Control 12 (2013) 351–358

Fig. 7. The thermodynamic properties of the reactions studied in this paper cal-culated from HSC Chemistry database and ab initio thermodynamic calculation:(t

teespIaPpwwiccTwhacapspr

mixing into the absorbent. The capacity was increased from ∼7 wt.%to 12 wt.%, also shown in Fig. 11. This confirms that the amount ofNaNO3 in the absorbent has a significant influence on the absorbent,and the capacity appears to correlate with NaNO3 loading.

a) enthalpy change versus temperatures and (b) Gibbs free energy change versusemperatures.

emperature and pressure is plotted in Fig. 8. According to thequilibrium curve, the reaction of Na2Mg(CO3)2 formation reachesquilibrium at PCO2 = 0.06 bar at 400 ◦C. Higher CO2 partial pres-ure is favorable for Na2Mg(CO3)2 formation, while at CO2 partialressures below 0.06 bar, Na2Mg(CO3)2 decomposition is favored.

n order to verify experimentally the equilibrium CO2 pressurest 400 ◦C, the Na–Mg double salt absorbent was evaluated throughSA, with the same test conditions described above, at different CO2artial pressures. At 400 ◦C, an equilibrium CO2 pressure of 0.4 baras observed, i.e., partial pressures of CO2 greater than 0.4 atmere required for CO2 absorption to occur. The experimental point

s shown in Fig. 8, and by applying the equation for equilibriumonstant K, ln K = −�G/RT (K equates PCO2), the predicted trendonsistent with that point is plotted and shown as a dashed line.his is higher than the theoretically derived value, but is consistentith the fact that the predicted equilibrium temperature is alsoigher than our experimentally derived value. Fig. 9 shows thatt a CO2 partial pressure of approximately 0.8 atm, the maximumapacity is reached within the 60 min allotted for the adsorption,nd a higher partial pressure of CO2 does not increase uptake. Inre-combustion, the warm syngas stream typically has a total pres-

ure of 15–20 bar with ∼20% CO2 content, which equals to a CO2artial pressure of 3–4 bar. With the capability of the absorbent toemove the CO2 partial pressure down to 0.4 bar, approximately

Fig. 8. CO2 partial pressure in equilibrium with the Na2Mg(CO3)2 double salt as afunction of temperature as predicted from ab initio calculation and after experimen-tal correction.

85–90% of the contained CO2 in the gas stream can be removed at400 ◦C through PSA. Similar results can also be achieved throughTSA.

3.4. Role of NaNO3

As discussed above, the absorption of CO2 is through theformation of Na2Mg(CO3)2 and it would appear that NaNO3 is sim-ply an impurity present in the absorbent. With careful control,for the absorbents with capacities of (15 ± 2) wt.%, prepared bythe described technique, the NaNO3 concentration is on average10–15 wt.%. Its selective removal would be expected to increase thepercentage of active components of the absorbent and the overallcapacity. During the preparation of the absorbent, after the dry-ing step, the contained Na2CO3 is fixed in the absorbent mixturethrough mostly forming insoluble Na2Mg(CO3)2, and some of theNaNO3 remaining with the absorbent can be removed through rins-ing with DI water over a filter. The washed powder was driedagain and activated at 400 ◦C for 3 h. The XRD pattern for theobtained absorbent in Fig. 10 exhibits no NaNO3 peak, indicatingthat the majority of the NaNO3 is removed. The absorbent washedto remove some NaNO3 was evaluated for CO2 capacity by TGAmeasurement. The absorbent without washing provided a baseline.The CO2 absorption capacity dropped by 50% after it was washed asshown in Fig. 11. 15 wt.% NaNO3 was then re-introduced by physical

Fig. 9. Experimentally measured data of the pressure dependence of Na2CO3-promoted MgO absorbent at 400 ◦C.

K. Zhang et al. / International Journal of Green

Fig. 10. X-ray diffraction patterns of the Na CO -promoted MgO absorbent washedt

Ib

oea1tcLtaiaN(adCt

anscrns

Fi

2 3

o reduce NaNO3.

t appears that NaNO3 plays an important role in the absorbenty facilitating the formation of Na2Mg(CO3)2.

NaNO3 melts and forms a thin layer of molten salt during theperation of both TSA and PSA processes. The effect of molten salt innhancing CO2 absorption with a solid absorbent at warm temper-ture was also observed in Li2ZrO3 powder (Ohashi and Nakagawa,999). The added K2CO3 was able to form eutectic molten salt withhe product compound Li2CO3. During continuous reaction, CO2ould diffuse through the liquid layer instead of through the solidi2CO3 particles as was the case without K2CO3. The differences inhe CO2 diffusion processes in the solid- and liquid phases presum-bly influence the CO2 absorption rate. However little informations known about the absorption product Na2Mg(CO3)2 and its inter-ction with NaNO3. Neither Na2CO3 nor MgCO3 form a eutectic withaNO3, although Na2CO3 has 3–4% dissolution in NaNO3 at 400 ◦C

FactSage, 2012). With its high wetting property over many oxidesnd metal surfaces, it is possible that NaNO3 penetrates the productouble salt grain boundaries and provides a liquid “channel” for fastO2 diffusion during absorption. Further study is clearly needed onhis point.

The observation of the important role of NaNO3 also points outnother reason for the difficulty of reproducing this synthesis tech-ique. During filtration, NaNO3 is retained in the absorbent in aimilar way as Na2CO3, which makes its final concentration poorlyontrolled and variable from batch to batch. In order to provide aeproducible and scalable synthesis, alternate synthesis methodseed to be developed. Progress on this topic will be reported in a

ubsequent publication.

ig. 11. TGA test results of the Na2CO3-promoted MgO absorbents showing thenfluence of NaNO3.

house Gas Control 12 (2013) 351–358 357

4. Conclusions

Experiments conducted with laboratory-synthesized Na2CO3-promoted MgO sorbents show a high CO2 capture capacity of3.4 mmol CO2/g sorbent in multiple cycle tests. These materialshold promise as CO2 absorbents for application from pre-combustion sources, in the warm temperature range 300–470 ◦C.The activated and regenerated absorbent comprises MgO, Na2CO3,and NaNO3, with MgO and Na2CO3 in proximity to facilitateNa2Mg(CO3)2 double salt formation on exposure to CO2. With theformation of the Na–Mg double salt, the temperature of operationof the sorbent is raised to the warm capture region of 300–470 ◦C.The heat of formation of the Na–Mg double salt, relative to its sep-arate component carbonates, has been determined through bothexperiment and computational calculation. This work has beensupported by thermodynamic analysis using HSC as well as withab initio thermodynamic calculations. Based on the equilibriumCO2 pressure at 400 ◦C for the double salt formation, the absorbentis predicted to be able to remove 85–90% of the CO2 in pre-combustion applications. The impurity phase, NaNO3, contained inthe absorbent is found to have an important role in enhancing theCO2 absorption. The challenges that have been encountered by oth-ers in reproducing the synthesis is due to the difficulty to controlthe amounts of the key components Na2CO3 and NaNO3, which arewater soluble and may vary from batch to batch. A modified or alter-nate synthesis technique needs to be developed in order improvematerial reproducibility and to allow scale up to larger quantitiesof material.

Acknowledgements

Financial support from the US DOE Office of Fossil Energy (NETL),the US DOE (EERE) Office of Biomass, the State of Wyoming, andPNNL internal investment (LDRD-ECI) is gratefully acknowledged.Some work was carried out at the Environmental and MolecularScience Laboratory, a national scientific user facility sponsored bythe DOE’s Office of Biological and Environmental Research (BER).

References

Bauer, T., Laing, D., Tamme, R., 2011. Recent progress in alkali nitrate–nitrite devel-opments for solar thermal power applications. In: Molten Salts Chemistry andTechnology Proceedings.

Blamey, J., Anthony, E.J., Wang, J., Fennell, P.S., 2010. The calcium looping cyclefor large-scale CO2 capture. Progress in Energy and Combustion Science 36,260–279.

Duan, Y., Sorescu, D.C., 2010. CO2 capture properties of alkaline earth metal oxidesand hydroxides: A combined density functional theory and lattice phonondynamics study. The Journal of Chemical Physics 133, 074508.

Duan, Y., Zhang, B., Sorescu, D.C., Johnson, J.K., 2011. CO2 capture properties ofM–C–O–H (M = Li, Na, K) systems: A combined density functional theory andlattice phonon dynamics study. Journal of Solid State Chemistry 184, 304–311.

FactSage, 2012. Na2CO3-NaNO3. In: Fact-Web Programs.Freeman, E.S., 1956. The kinetics of the thermal decomposition of sodium nitrate and

of the reaction between sodium nitrite and oxygen. Journal of Physical Chemistry60, 1487–1493.

Gregg, S.J., Ramsay, J.D., 1970. Adsorption of carbon dioxide by magnesia studied byuse of infrared and isotherm measurements. Journal of the Chemical Society A:Inorganic, Physical, and Theoretical, 2784.

Hassanzadeh, A., Abbasian, J., 2010. Regenerable MgO-based sorbents for high-temperature CO2 removal from syngas: 1, Sorbent development, evaluation, andreaction modeling. Fuel 89, 1287–1297.

Herzog, H., 2009. Carbon dioxide capture and storage. In: Helm, D., Hepburn, C.(Eds.), The Economics and Politics of Climate Change. Oxford University Press,pp. 263–283.

Hufton, J.R., Waldron, W., Weigel, S.J., Rao, M., Nataraj, S., Sircar, S., 2000. Sorptionenhanced reaction process (SERP) for production of hydrogen. In: Proceedingsof the 2000 Hydrogen Program Review.

Mayorga, S.G., Weigel, S.J., Gaffney, T.R., Brzozowski, J.R., 2001. Carbon dioxide adsor-bents containing magnesium oxide suitable for use at high temperatures. US6,280,503 B1, Air Products and Chemical, Inc., United States.

Nissen, D.A., Meeker, D.E., 1983. Nitrate/nitrite chemistry in sodium nitrate–potassium nitrate melts. American Chemical Society 22, 716–721.

3 f Green

O

S

S

58 K. Zhang et al. / International Journal o

hashi, T., Nakagawa, K., 1999. Effect of potassium carbonate additive on CO2

absorption in lithium zirconate powder. Materials Research Society 547,249–254.

ingh, R., Ram Reddy, M.K., Wilson, S., Joshi, K., Diniz da Costa, J.C., Webley,

P., 2009. High temperature materials for CO2 capture. Energy Procedia 1,623–630.

ircar, S., Anand, M., Carvill, B., Hufton, J.R., Mayorga, S.G., Miller, B., 1995. Sorptionenhanced reaction process (SERP) for production of hydrogen. In: Proceedingsof the 1995 U.S. DOE Hydrogen Program Review.

house Gas Control 12 (2013) 351–358

Siriwardane, R.V., Stevens, R.W., 2009. Novel regenerable magnesium hydroxidesorbents for CO2 capture at warm gas temperatures. Industrial and EngineeringChemistry Research 48, 2135–2141.

Steeneveldt, R., Berger, B., Torp, T., 2006. CO2 capture and storage closing

the knowing–doing gap. Chemical Engineering Research and Design 84,739–763.

Xiao, G., Singh, R., Chaffee, A., Webley, P., 2011. Advanced adsorbents based on MgOand K2CO3 for capture of CO2 at elevated temperatures. International Journal ofGreenhouse Gas Control 5, 634–639.