Journal Identification = IJGGC Article Identification = 418 Date: July 6, 2011 Time: 2:50 pm

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International Journal of Greenhouse Gas Control 5 (2011) 634–639

Contents lists available at ScienceDirect

International Journal of Greenhouse Gas Control

j our na l ho me p age: www.elsev ier .com/ locate / i jggc

dvanced adsorbents based on MgO and K2CO3 for capture of CO2 at elevatedemperatures

ongkui Xiaoa, Ranjeet Singha, Alan Chaffeeb, Paul Webleya,∗

Cooperative Research Center for Greenhouse Gas Technologies, Department of Chemical Engineering, Monash University, Wellington Road, Clayton, Victoria, 3800, AustraliaCooperative Research Center for Greenhouse Gas Technologies, School of Chemistry, Monash University, Wellington Road, Clayton, Victoria, 3800, Australia

r t i c l e i n f o

rticle history:eceived 3 October 2010eceived in revised form 10 April 2011ccepted 12 April 2011vailable online 18 May 2011

eywords:O2 capture

a b s t r a c t

We report here our study on development of adsorbents suitable for capturing CO2 from synthesis gas(syngas) at high temperatures (>100 ◦C). Our adsorbents are based on double salts of MgCO3 and K2CO3

in which samples with different ratios of Mg:K were prepared by the wet mixing method of magnesiumnitrate and potassium carbonate. The adsorbents were characterized by X-ray diffraction analysis, ther-mogravimetric analysis (TGA), and N2 adsorption and desorption at 77 K. The morphology of the sampleswas observed with scanning electron microscopy (SEM). CO2 adsorption experiments were performedat four temperatures, 300 ◦C, 350 ◦C, 375 ◦C, and 400 ◦C, which correspond to the operating temperature

dsorptionre-combustionigh temperature

range of gases exiting a typical water gas shift reactor and entering a gas turbine in the Integrated Gasi-fication Combined Cycle (IGCC) process. The CO2 adsorption amounts were 1.65 wt%, 8.47 wt%, 8.55 wt%,and 0.63 wt% respectively at a CO2 partial pressure of 100 kPa. These adsorbents also exhibited excellentcycling stability both in temperature swing adsorption and pressure swing adsorption operation. Thekinetics of CO2 adsorption at different temperatures was obtained by using the linear driving force (LDF)model and a diffusion activation barrier of 21 kJ/mol was also inferred.

. Introduction

Global warming is attributed to the accumulation of carbonioxide (CO2) in the atmosphere from anthropogenic sources. Flueas from coal-fired power plants is one of the major sources ofO2 emissions (Schrag, 2007). Therefore, there is an urgent need toevelop and deploy technologies for CO2 removal from flue gasesesulting from coal burning. IGCC is a promising technology foroal-fired power plants and provides an opportunity for capturingO2 prior to combustion of the syngas (Shinada et al., 2002), thusroducing a CO2-free emission from the turbine. This CO2 separa-ion process is referred to as pre-combustion CO2 capture. In IGCCrocesses, the gases exiting the gasifier after gas cleanup and ini-ial cooling is usually around 400 ◦C and the gases entering the gasurbine (often after a water-gas shift reactor) is typically 250 ◦C. As

result, the best temperature window for separating CO2 woulde in the range of 250–400 ◦C (Siriwardane et al., 2007). Various

ethods for pre-combustion CO2 capture have been reported, such

s absorption with solvents (Hagewiesche et al., 1995; Maceirast al., 2008), separation with membranes (Amelio et al., 2007;

∗ Corresponding author. Tel.: +61 3 99053445, fax: +61 3 99055686.E-mail addresses: james.xiao@monash.edu (G. Xiao), ranjeet.singh@monash.edu

R. Singh), alan.chaffee@monash.edu (A. Chaffee), paul.webley@monash.eduP. Webley).

750-5836/$ – see front matter © 2011 Elsevier Ltd. All rights reserved.oi:10.1016/j.ijggc.2011.04.002

© 2011 Elsevier Ltd. All rights reserved.

Mavroudi et al., 2003; Paolo et al., 2007), and adsorption with solidsorbents (Lee et al., 2008a,b). In spite of the fact that absorptionwith solvents is currently the most widely used technology forCO2 capture, this process occurs at relatively low temperaturesand therefore requires cooling of the syngas followed by reheat-ing prior to admission to the turbine. This cooling and reheating isenergy inefficient although part of the cooling is usually recovered(Ciferno et al., 2008). By contrast, membrane- and adsorbent-basedprocesses have the potential for CO2 capture at higher temperatureswith elimination of the intermediate cooling steps.

The most important aspect of an adsorption-based process is theselection of the adsorbent(s). Several materials have been devel-oped in recent years which selectively and reversibly adsorb CO2at elevated temperatures (Lee et al., 2008a). Among them, cal-cium oxide related materials (Abanades, 2002; Abanades et al.,2005; Grasa and Abanades, 2006; Manovic and Anthony, 2009;Martavaltzi and Lemonidou, 2008; Reddy and Smirniotis, 2004)and hydrotalcite-like compounds (Hutson and Attwood, 2008; Leeet al., 2007a,b; Oliveira et al., 2008; Reynolds et al., 2006; Wanget al., 2008) are the most widely reported adsorbents. Calcium oxidecaptures CO2 by the carbonation reaction at around 600 ◦C andthe captured CO2 can be desorbed at temperatures above 800 ◦C,

thus leading to the regeneration of the adsorbent for reuse. MostCaO related materials such as dolomite and limestone are read-ily available in nature. However, regeneration of these materialsat high temperatures is always accompanied by a rapid sintering

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Greenhouse Gas Control 5 (2011) 634–639 635

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G. Xiao et al. / International Journal of

rocess. The naturally available and man-made hydrotalcite mate-ials have been reported to be suitable for in situ CO2 separationt high temperatures in a pressure swing adsorption process inater gas shift reactors, but the adsorption amount is relatively

ow (e.g., 0.65 mol/kg at 400 ◦C) (Ding and Alpay, 2000). The mech-nism of CO2 adsorption on hydrotalcite is not clear, though severaleports have investigated this subject (Ebner et al., 2006; Hutsont al., 2004; Pérez-Ramírez et al., 2007; Walspurger et al., 2008). Aetailed comparison of adsorption performance of widely reporteddsorbents can be found elsewhere (Lee et al., 2008a,b; Yong et al.,002).

Alkali metal promoted metal oxides have also received increas-ng attention in recent years. Chemisorption isotherms of Na2Oromoted alumina have been reported by Lee et al. (2007a,b) forO2 capture at temperatures of 250, 350 and 450 ◦C, and the adsorp-ion capacities are approximately 1 mol/kg. Potassium carbonateupported on MgO (Lee et al., 2008a,b) has also been reported toe capable of capturing CO2 at elevated temperatures with goodycling stability. The adsorption capacities of these materials aretill relatively low for large scale usage; therefore, it is necessary toevelop some adsorbents with higher CO2 adsorption capacities.

Recently, Mayorga et al. (2001) reported novel magnesium oxidedsorbents with potential for CO2 capture. Some of the materi-ls reported in these patent exhibited remarkable CO2 capacitiest high temperatures but no detailed account was presented. Ourork reported here was motivated by this report and we have

xamined these materials and modifications thereof in more detailo understand their potential application. This adsorbent is suitableor CO2 adsorption in the temperature range of 250–400 ◦C, whichs the typical temperature exiting a coal gasifier in an IGCC pro-ess. In this study, physical properties of the prepared adsorbentsuch as surface area and structures before/after CO2 adsorptionere investigated by gas adsorption analysis and powder XRD.

he morphology of the adsorbents was observed by using scan-ing electron microscopy. The CO2 adsorption performance was

nvestigated with the aid of a thermogravimetric analyser, and thedsorption kinetics were obtained and represented using the linearriving force model.

. 2 Experimental work

.1. Chemicals and gases

In this study, magnesium nitrate hexahydrate and potassiumarbonate anhydrous (both from Ajax Finechem, 99%) were usedor the preparation of the MgO based adsorbents. Nitrogen, used asurge gas during calcination and the cyclic desorption period, had

purity of 99.999%. The purity of carbon dioxide used for adsorp-ion was 99.995%. All chemicals were used as received without anyurther purification. Deionised water was used in all experiments.

.2. Sorbent synthesis

MgO based adsorbents were prepared with different startingolar ratios of magnesium nitrate hexahydrate (Mg(NO3)2·6H2O)

o potassium carbonate (K2CO3) according to a similar procedures that reported by Mayorga et al. (2001). A specified amount ofagnesium nitrate hexahydrate was first dissolved in 500 mL of

eionised water, and the corresponding desired molar amount ofotassium carbonate powder was quickly added to the previouslyormed solution under vigorous magnetic stirring. Upon mixing, the

olution turned to dilute slurry. Stirring of the slurry was continuedntil there was no apparent remaining free water. The sticky slurryas kept at 25 ◦C overnight in open air and filtered to certain dry-ess. The filter cake was then dried at 120 ◦C for 24 h in oven. The

Sample DS1 DS2 DS3 DS4

Mg:K 1:0.45 1:1.89 1:2.13 1:1.84

as-prepared samples were then calcined in a tube furnace at 450 ◦Cfor 3 h in nitrogen flow to obtain the corresponding sorbents. Theratios of Mg(NO3)2·6H2O to K2CO3 for the four adsorbents were 1:2,1:3, 1:4 and 1:5, correspondingly, the resultant prepared samplesare designated as DS1, DS2, DS3 and DS4, respectively.

2.3. Characterization

Powder X-ray diffraction (XRD) was recorded with a PhillipsAutomated Powder Diffractometer with Cu-K� radiation (40 kV,25 mA) at a wavelength of 1.54 A. Thermogravimetric (TG) analy-ses were conducted using a Mettler Toledo TGA-SDTA851 analyser(Switzerland) from 25 to 800 ◦C under nitrogen (20 mL/min) toinvestigate thermal decomposition of the adsorbent. Nitrogenadsorption/desorption isotherms were measured at 77 K withan ASAP 2020 analyser (Micromeritics, USA), and the sampleswere degassed at 350 ◦C for 12 h prior to each measurement. TheBrunauer–Emmett–Teller (BET) method was adopted to calculatethe specific surface areas. Morphology of the samples was obtainedwith JEOL 6300F FEGSEM microscope (Japan) operated at 15 kV.

2.4. CO2 adsorption measurement

Gravimetric determination of carbon dioxide uptake at 300, 350,375 and 400 ◦C at atmospheric pressure were performed using athermogravimetric analyser (Mettler Toledo TGA-SDTA851). Thecalcined adsorbents were first activated at 450 ◦C with a temper-ature ramp rate of 10 ◦C/min and kept at 450 ◦C for 60 min undernitrogen flow before temperature was reduced to the correspond-ing adsorption temperatures at a rate of −5 ◦C/min and kept atthe adsorption temperature for another 15 min with nitrogen flow.Then the gas was switched to CO2 for the adsorption phase. Theadsorption time was set to be 120 min when a near equilibriumadsorption state could be obtained.

3. Results and discussion

3.1. Structural and textual properties

Wide-angle XRD patterns of the prepared samples were col-lected to analyse phase changes in the samples before and afterCO2 adsorption. As seen in Fig. 1a, all of the four freshly calcinedsamples before CO2 adsorption show two phases: one is MgO(JCPDS 870653) and the other is K2CO3 (JCPDS 711466). However,the relative strength of MgO peaks to K2CO3 peaks are different,indicating the ratios of MgO to K2CO3 are different in each sam-ple which could be further demonstrated by elemental analysis(Table 1). The MgO in the samples resulted from the decomposi-tion of MgCO3 in the precipitate at the calcination temperatureof 450 ◦C and K2CO3 was the stoichiometric excess K2CO3 in theprecipitate as it did not decompose at 450 ◦C (Stern, 2000). XRDpatterns of the samples after CO2 adsorption (Fig. 1b) illustratethree phases, which are MgO, K2CO3, and a new phase which can beindexed as K2Mg(CO3)2 (JCPDS 751725). The newly formed species,K2Mg(CO3)2, also reported by Lee et al. (2008a,b) for low tem-perature CO2 adsorption in the presence of water vapour using

MgO based adsorbents, was the reaction product of MgO, K2CO3with CO2. There were no KHCO3 or MgCO3 phases found after CO2adsorption, in contrast to the results of CO2 adsorption at 60 ◦Creported by Lee et al. (2008a,b), as their respective decomposition

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3.2.2. Adsorption at different temperatures

Fig. 1. Powder XRD patterns of the samples: (a) before CO2 adsor

emperatures are 100 ◦C to 120 ◦C and 330 ◦C, much lower than thedsorption temperature.

N2 adsorption/desorption isotherms show that the as-prepareddsorbents have a hysteresis loop at p/p0 of 0.8–1.0, indicating largeoid spaces between crystals (Fig. 2). The data calculated by theET method show that surface areas of the prepared samples areround 2 m2/g, which is much lower than most of the adsorbents forhysisorption and essentially corresponds to the external surfacerea of the crystals.

Differential thermogravimetric curves of the four samples wereecorded over the temperature range of 25–800 ◦C in nitrogenow (Fig. 3). These curves show that the highest weight loss rateccurred at different temperatures, increasing with the contentf K2CO3 in the starting material (Fig. 3). For DS1, the highestecomposition rate appeared at 331 ◦C, which is the same decom-osition temperature as pure MgCO3 (Stern, 2000). This is expecteds MgCO3 is the main phase in this sample since the molar ratiof MgCO3 to K2CO3 is 1:0.45. For samples with higher potassiumontent, the corresponding highest rate of decomposition occurredt 388 ◦C for DS2, 423 ◦C for DS3, and 489 ◦C for DS4. The shift ofighest weight loss peaks from normal decomposition temperaturef MgCO3 to higher temperature is in keeping with the relativencrease in K2CO3 and its corresponding higher decompositionemperature (K2CO3).

The morphology of the prepared adsorbents as seen by their

EM images is shown in Fig. 4. It is clear that the adsorbents aregglomerates of small crystals, and the particle size increased withhe increase in K2CO3 content.

ig. 2. N2 adsorption-desorption isotherms at 77 K for the adsorbent samples.

(b) after CO2 adsorption: (�) MgO; (©) K2CO3; (∀) K2Mg(CO3)2.

3.2. CO2 sorption experimental results

3.2.1. Influence of Mg/K ratio of the starting materialsThe influence of the ratio of K to Mg of the starting materials on

CO2 adsorption capacities of the prepared samples was examinedby performing CO2 (99.995%) adsorption experiments on a thermo-gravimetric analyser (TGA) at 375 ◦C (Fig. 5). It is clear from Fig. 5that adsorbents with different starting material ratios have differ-ent adsorption capacities. At a ratio of 1:2 (DS1), the adsorptionamount was about 5.28 wt% increasing to 8.55 wt% when the ratioswere increased to 1:3 (DS2). Further increase to 1:4 (DS3), and 1:5(DS4) did not lead to additional CO2 adsorption capacity. The ratioof 1:3 was preferred as this minimizes usage of starting materi-als. The effect of K:Mg ratio on CO2 adsorption capacities could beexplained from XRD analysis and elemental analysis. XRD patternsshow that the adsorption of CO2 on these samples resulted from thereaction between MgO, K2CO3, and CO2, at the same time, the ele-mental analysis shows that molar ratios of MgO to K2CO3 in DS1,DS2, and DS3 are 1:0.45, 1:0.95, 1:1.06, and 1:0.92, respectively.In DS2, DS3, and DS4 the ratios were close to the stoichiometricratio of MgO/K2CO3 in the reaction, therefore there was no excessof either component which could take up some dead weight of thesample as was the case for DS1.

Our initial work suggested a ratio of 1:3 (sample DS2) was suf-ficient to exhibit good adsorption capacity without excessive use

Fig. 3. Differential thermogravimetric profiles of adsorbent samples.

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of CO2 at 400 ◦C was also verified on the TGA. As 375 ◦C is a suit-

Fig. 4. SEM images of the p

f K2CO3 and hence we chose this sample for further analysis. Thedsorption of CO2 at different temperatures was performed on sam-le DS2. The specific temperatures were chosen because of theppropriate temperature range for CO2 capture in a typical IGCClant as discussed earlier. Fig. 6 shows that the weight change ofhe adsorbent at 300 ◦C is a gradual process increasing slowly with

final CO2 uptake of 1.2 wt% of the adsorbent (Fig. 6). At 350 ◦Cnd 375 ◦C however, the weight of the adsorbent experienced aapid increase with final CO2 uptakes of 8.39 wt% and 8.69 wt%

or 350 and 375 ◦C, respectively. Adsorption of CO2 at 400 ◦C wasegligible. The variation of adsorption amount with temperaturelso verifies that the adsorption of CO2 on these materials is

Fig. 5. Effect of K:Mg ratio on CO2 adsorption capacity.

ed double salt adsorbents.

chemisorption and there is an optimal temperature for the chem-ical reaction. As is well known, physisorption mechanism leads toa decrease in the amount adsorbed with an increase of tempera-ture. The adsorption capacities at 350 and 375 ◦C are much higherthan those of widely reported hydrotalcite-like adsorbents whichwere tested for their use in water gas shift reactors (Wang et al.,2008). Therefore, this material has great potential for high tem-perature CO2 adsorption in IGCC processes. The lack of adsorption

able adsorption temperature for this adsorbent, this temperaturewas adopted as the adsorption temperature for further adsorptionexperiments.

Fig. 6. Uptake curves of the adsorption of pure CO2 at different temperatures onsample DS2.

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ure of 1 bar: (a) pressure swing, (b) temperature swing with nitrogen purge.

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Fig. 7. Cyclic experiments of CO2 adsorption at 375 ◦C with a CO2 press

.2.3. Cycling stabilityGood adsorbents should retain their cyclic adsorption capacity

uring the adsorption process. Two sets of CO2 adsorp-ion/desorption cycling experiments were conducted on thedsorbent with a Mg:K ratio of 1:3. In the first set (pressure swingdsorption (PSA)) CO2 adsorption and desorption by N2 purge bothccurred at 375 ◦C. In the second set (temperature swing adsorp-ion (TSA)) CO2 adsorption took place at 375 ◦C and desorption at00 ◦C under N2 purge. For both of the two experiments, the adsorp-ion time was 20 min and desorption time was 30 min. 17 cyclesere carried out for the PSA (Fig. 7a) experiment and 9 cycles for

he TSA experiment (Fig. 7b). For the PSA experiment, there was aontinuous baseline drift throughout the 17 cycles, but the adsor-ent showed stable cycling adsorption capacity, and did not losedsorption capacity. After 17 cycles the sacrificed capacity percent-ge was only about 5.1% of the original capacity (from 7.9 wt% to.5 wt%). The phenomena of baseline drift and almost stable adsorp-ion capacity may be explained on the basis that the adsorbed CO2as only partially regenerated in each cycle because of insufficient

ime for desorption and the un-regenerated part did not hinder fur-her adsorption of CO2. For the TSA experiment, as the adsorbentould be fully regenerated by increasing the temperature to 400 ◦C,he curve showed a flat baseline and the adsorption capacity waslmost the same in each of the 9 cycles. More cyclic stability testsver 100 s of cycles are needed to confirm long term stability.

.2.4. Kinetics of adsorption of CO2As shown in Fig. 6 for the CO2 adsorption on DS2, the rate of

dsorption was initially fast, and then slowed to a nearly steadytate at 375 ◦C and 350 ◦C, while for the adsorption at 300 ◦C, thentire uptake adsorption process showed a gradual increase. TheO2 uptake may be considered to be a pseudo-first-order massransfer between the gas phase and the adsorbent for the first0 min at which point the adsorption amount had virtually reachedquilibrium. The linear driving force (LDF) model (also equivalent to

first order reaction model mathematically) was used to calculatehe uptake rate constant for the initial stage.

n[

1 − Mt

Me

]= kt (1)

here k is the rate constant, Mt is the gas uptake at time t, Me is theas uptake at equilibrium, and t is the adsorption time. The plot ofn(1 − Mt/Me) against time t is shown in Fig. 8.

The adsorption kinetics at different temperatures were com-

ared in terms of the rate constant (k) of CO2 adsorption as shown

n Fig. 8. The data show good fit to Eq. (1) at 375 ◦C with deterio-ating agreement at 350 and 300 ◦C. The adsorption rate constantsbtained from the slope of the linear fit for 375 ◦C, 350 ◦C and 300 ◦C

Fig. 8. Kinetics of CO2 adsorption on adsorbent DS2.

are 0.089 min−1, 0.078 min−1, and 0.056 min−1, respectively. Thedecrease in rate constant from high to low temperature is fur-ther evidence that the adsorption of CO2 on the adsorbent was achemisorption process. The activation energy for this rate constantis 21 kJ/mol which is relatively high for modest diffusion barrierand may be indicative of an impermeable shell of reacted adsorbentcoating the material as the reaction proceeds. This activation energyis lower than the activation energy (37.5 kJ/mol) reported by Leeet al. (2007a,b) on the sodium oxide promoted alumina adsorbents.

4. 4 Conclusions

This work has studied some adsorbents suitable for capturingCO2 from syngas at high temperatures. Adsorbents with differentratios of Mg/K have been prepared by the wet mixing method ofmagnesium nitrate and potassium carbonate. The adsorbents havebeen characterized by X-ray diffraction analysis, thermogravimet-ric analysis (TGA), and N2 adsorption and desorption at 77 K. It wasfound that there were only MgO and K2CO3 phases presented in theadsorbents before CO2 adsorption, and a new phase formed afterCO2 adsorption. The surface area of these adsorbents was very low(∼10 m2/g) compared to most physical adsorbents though similar

to that of some chemisorbents. CO2 adsorption experiments havebeen performed at three different temperatures, 300 ◦C, 350 ◦C, and375 ◦C, which are the temperature window between the gases exit-ing a typical water gas shift reactor and entering a gas turbine.

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he corresponding adsorption amounts are 2.8 wt%, 6.5 wt%, and.99 wt% of the sample weight. These adsorbents have also shownxcellent cycling stability both in temperature swing adsorptionnd pressure swing adsorption. They did not experience the samessue of deterioration as the widely reported high temperaturedsorbents in literature. The kinetics of CO2 adsorption at differ-nt temperatures was obtained by using the LDF model and wasound to be 0.089 min−1 at 375 ◦C. A diffusion activation barrier of1 kJ/mol was found.

cknowledgements

We acknowledge the funding provided by the Australian Gov-rnment through its CRC Program to support this CO2CRC researchroject. The authors appreciate the helpful discussion with Zhangx-

ong Wu, and the assistance from staff in Monash Centre forlectron Microscopy.

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