Synthesis of efficient CaO sorbents for CO2 capture using a simple

The effects of the decomposition atmosphere, carbon source. ... sintering of CaCO3, and the subsequent carbon burn-off in air promoted the fabrication...
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Synthesis of Efficient CaO Sorbents for CO2 Capture Using a Simple Organometallic Calcium-Based Carbon Template Route Pengfei Zhao, Ji Sun, Yimin Li,* Ke Wang,* Zeguang Yin, Zhongyun Zhou, and Zhen Su School of Electric Power Engineering, China University of Mining and Technology, Xuzhou, Jiangsu 221116, People’s Republic of China S Supporting Information *

ABSTRACT: A facile carbon template method was employed to modify the microstructure of calcined organometallic calcium compounds (i.e., calcium acetate, calcium citrate, and calcium gluconate) for high-temperature CO2 capture. The effects of the decomposition atmosphere, carbon source, and pyrolysis temperature on the physical and chemical properties as well as the cyclic CO2 capture performance were determined using various morphological characterization techniques and detailed thermogravimetric analysis. During pyrolysis in an inert atmosphere, in situ carbon formed from organometallic calcium compounds, which served as a template for controlling the sintering of CaCO3, and the subsequent carbon burnoff in air promoted the fabrication of porous CaO. Among the three organometallic calcium compounds, the pyrolysis process of calcium gluconate exhibited the earliest carbonization and the template with the highest organic carbon content and the slowest decomposition of CaCO3. Therefore, the most favorable structure (i.e., the smallest crystal size, largest specific surface area, and largest pore volume of CaO) was obtained from calcium gluconate, which was responsible for the fastest adsorption rate, highest capacity, and best stability. Moreover, this superior performance was maintained when the pyrolysis temperature was approximately 600−800 °C.

1. INTRODUCTION Anthropogenic CO2 emissions from fossil fuel combustion power plants and other industrial processes are most likely the main driving force in global warming and environmental change.1 CO2 capture and storage (CCS) are important solutions for mitigating climate change.2,3 Among the current CO2 capture technologies, calcium looping based on cyclic carbonation/decarbonation of CaO has recently attracted much attention as a result of the low costs of naturally available CaO precursors (limestone) and the high theoretical capture capacity of 0.78 g of CO2/g of CaO.4,5 However, limestone-derived CaO suffers from a major drawback [i.e., the rapidly decreasing CO2 uptake capacity after several cycles as a result of the low Tammann temperature of CaCO3 (∼529 °C)].6,7 Thermal sintering induces severe aggregation of crystals and a dramatic reduction in their specific surface areas and pore volumes.8 These morphological changes shift the CO2 capture mechanism from the rapid chemically controlled regime to the slow diffusioncontrolled regime.9 Therefore, an irreversible loss of CaO reactivity occurs within a given residence time in the reactor.10 To improve the sintering resistance and mitigate the decrease in the CO2 uptake capacity, numerous methods, such as the identification of potential sintering-resistant calcium precursors,11−14 steam reactivation/pretreatments,15−18 and the incorporation of refractory materials,19−27 have been proposed. Calcium precursors are important for developing a high and stable CO2 uptake capacity in CaO sorbents. Lysikov et al.28 observed that asymptotic CO2 uptake characteristics were affected by the morphology of the CaO precursor (e.g., surface area, pore volume, and pore size distribution). In general, a tailored CaO sorbent with a large surface area and pore volume results in high carbonation conversion. An emerging route © 2016 American Chemical Society

involving the direct calcination of organometallic precursors in air can produce CaO with a high surface area. Lu et al.29 determined that CaO derived from Ca(CH3COO)2 possessed a higher CO2 uptake capacity than CaO derived from other calcium precursors. Liu et al.30,31 further demonstrated that the decomposition of most of the organometallic precursors resulted in a larger specific surface area and porosity. In particular, CaO derived from calcium gluconate exhibited the highest CO2 uptake capacity. Alternatively, Li et al.32 reported the acidification of natural limestone using an acetic acid solution to improve the durability and uptake capacity of natural limestone. Other organic acids, including formic acid33 and tartaric acid,34 have also been employed to enhance the cyclic adsorption capacity of limestone. However, these organic calcium compounds or acetified limestone must be calcined at high temperatures (i.e., >800 °C) through a CaCO3 intermediate to produce CaO. Direct calcination in air induced a thermal sintering of CaCO3, which caused undesirable grain growth and agglomeration.35 To overcome this drawback, Armutlulu et al.36 developed a novel carbon-based sacrificial templating approach to tailor hollow nanostructured CaO sorbents. However, this method involved multistep procedures and expensive reagents. Recently, Radfarnia et al.37 introduced a facile two-step calcination of citricacid-treated limestone to prepare highly porous CaO. In their study, in situ carbon that was formed during the pyrolysis step under argon acted as a CaO dispersant during burnoff in the secondary calcination step under air. Wang et al.38 also applied Received: May 22, 2016 Revised: August 22, 2016 Published: August 24, 2016 7543

DOI: 10.1021/acs.energyfuels.6b01223 Energy Fuels 2016, 30, 7543−7550

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Figure 1. Pyrolysis diagrams for CG, CC, and CA under a nitrogen atmosphere and CG under an air atmosphere: (a) TG curves and (b) DTG curves. pattern was determined using the Barrett−Joyner−Halenda (BJH) method. The surface morphology was analyzed using a JSM-6360LV scanning electron microscope (SEM). The sample was spread directly on conductive tape followed by metal spraying preprocessing. Then, SEM was employed at different magnifications to observe the surface morphology of the sample. The transmission electron microscopy (TEM) images were captured on a JEM-2010. The crystalline structure was analyzed using a Rigaku D/max2550 X-ray diffractometer (XRD). The wide-angle scan parameters included a Cu Kα target (λ = 1.540 56 Å), a scan voltage of 40 kV, a current of 100 mA, a scan rate of 10°/min, a step size of 0.02°, and a scan range from 10° to 90°. The elemental carbon content was measured using an element analyzer (Germany Elementar Vario III). 2.3. Cyclic Adsorption Test. The cyclic CO2 performance was tested using a thermogravimetric analyzer (ZRY-1P, Techcomp Jingke Scientific Instrument Co., Ltd., Shanghai, China). For each run, approximately 5 mg of the sample was placed in an alumina crucible. Prior to the cycling tests, the sample was calcined (800 °C at 15 °C/ min) in a 40 mL/mg N2 flow for 5 min. Then, the temperature was decreased to 700 °C at a rate of 15 °C/min. The cyclic test began with carbonation (15 vol % CO2 in N2 for 20 min). After carbonation, the temperature was increased to 800 °C (15 °C/min) and the sorbent was regenerated in the same mixed gas (15 vol % CO2 in N2) for 5 min. This process was repeated for 10 carbonation/calcination cycles.

the citric-acid-based carbon template to synthesize an efficient MgO-stabilized CaO sorbent. The citric-acid-based carbon template is a promising method for the fabrication of structurally improved CaO sorbents. However, more effective organic-acidbased template methods may exist. Therefore, in this study, we extend the carbon template method using calcium gluconate as a carbon source to prepare a highly effective, pure CaO sorbent. More importantly, on the basis of the morphology and cyclic sorption capacity characteristics, the effects of the carbonization process, carbon source (calcium gluconate versus calcium acetate versus calcium citrate), and carbonization temperature were investigated to elucidate the carbon template mechanism. This technique relies on in situ carbonation of organic-acid-based calcium, which avoids the further addition of a carbon template during the synthesis.

2. EXPERIMENTAL SECTION 2.1. Sorbents. The analytical reagents used in this study included calcium gluconate monohydrate (CG), calcium citrate tetrahydrate (CC), and calcium acetate monohydrate (CA). All of the reagents were obtained from Shanghai Aladdin Reagent Co., China. The carbon template was prepared as follows: the proper amount of organic-acidbased calcium was placed in a corundum crucible and heated by a tube furnace under a nitrogen atmosphere at a heating rate of 10 °C/min to its carbonization temperature (600, 800, or 900 °C) to obtain a black powder. Then, the atmosphere was changed to normal air, and the material was calcined at 800 °C for 2 h to produce the CaO sorbents. To characterize the synthesis technique, the following nomenclature is used: The first part of the label (i.e., CG, CC, and CA) represents the calcium precursor (calcium gluconate, calcium citrate, and calcium acetate). In the second part of the label, 6, 8, and 9 represent the carbonization temperatures in N2 (i.e., 600, 800, and 900 °C). In the third part of the label, “a” represents further air calcination at 800 °C. For example, CG-6-a refers to calcium gluconate that has been carbonized at 600 °C in N2 followed by calcination in air. For comparison, the direct calcination method was employed as follows: the proper amount of the sample (calcium gluconate, calcium citrate, and calcium acetate) was calcined directly in an air atmosphere at 800 °C (using a ramp up of 10 °C/min) for 2 h. The resulting sorbents were named CG-a, CC-a, and CA-a. 2.2. Characterization Methods. A thermogravimetric analyzer (Mettler Toledo) was used in the decomposition test. Approximately 10 mg of the original sample was placed in an alumina crucible and heated from room temperature to 1000 °C at a rate of 10 °C/min under a nitrogen or air atmosphere with a flow rate of 50 mL/min. A Micromeritics ASAP-2010C nitrogen adsorption instrument was employed to analyze the texture of the sample. Prior to the test, approximately 2.5 g of original sample was subjected to vacuum degassing for at least 3 h, and its specific surface area was calculated using the Brunauer−Emmett−Teller (BET) method. In addition, its pore

3. RESULTS AND DISCUSSION 3.1. Effect of Decomposition Atmospheres. Figure 1 shows a diagram of the thermal decomposition [including the thermogravimetric (TG) and derivative thermogravimetric (DTG) curves] of CG under nitrogen (CG-N2) and air (CGair) atmospheres. The detailed weight loss curves are presented in Figure S1 of the Supporting Information. According to the TG

Figure 2. XRD patterns of CG-a and CG-6-a. 7544

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which was primarily due to the different amounts of carbon deposited. This result is similar to that reported by Labuschagne et al.39 In addition, the differential scanning calorimetry (DSC) tests of Labuschagne et al. revealed that CG-air exhibited an intense exothermic peak at approximately 550 °C, which suggests that deposited carbon burned violently. In contrast, CG-N2 showed a relatively mild exothermic peak. Therefore, the deposited carbon of CG-N2 could accumulate. On the basis of the initial organic content and the final weight, the final organic carbon amount of CG-N2 is ∼44.40%. The third loss (stage III), which was initiated at ∼600 °C, was mainly due to the calcination of CaCO3. Under an air atmosphere, this decomposition process occurred quickly and was complete at ∼680 °C. In contrast, as a result of deposited carbon, CG-N2 exhibited a slow decomposition. Furthermore, this decomposition shifted slightly to higher temperatures and was complete at ∼900 °C. The large differences in the decomposition process of CG resulted in the formation of CaO sorbents with different structures. Figure 2 shows the XRD patterns of CG-a and CG-6-a. Both samples had the same characteristic peaks of highly crystalline CaO. The average crystallite size was estimated using Scherrer’s equation: D (nm) = kλ/β cos θ, where k is the shape factor (k = 0.89), λ is the wavelength of X-rays, β is the full width at half maximum, and θ is the diffraction angle. The crystal size of CaO (Table 1) in CG-6-a (15.6 nm) was only half that of CaO in CG-a (32.7 nm). Additionally, both CC-6-a and CA-6-a gained CaO sorbents with smaller crystal sizes compared to CC-a and CA-a. Therefore, deposited carbon (stage II in Figure 1) acted as a physical barrier to separate the intermediate CaCO3 grains and effectively hinder both crystal growth and agglomeration at high temperatures (stage III in Figure 1). Figure 3 compares the SEM images of CG-a and CG-6-a. CG-a (Figure 3a) possessed a flaky structure with an uneven and relatively large foliation. However, CG-6-a (Figure 3b) exhibited a relatively uniform blade-like structure. The results from N2 adsorption (Table 2) further revealed that the specific surface area, pore volume, and average pore size of CG-6-a were relatively larger than those of CG-a. The cyclic CO2 capture characteristics of both CaO sorbents are shown in Figure 4. The maximum initial CO2 uptake capacity of CG-a was 0.605 g of CO2/g of sorbent (g/g). This result is in agreement with the results reported by Liu et al.31 Moreover, the maximum initial CO2 uptake capacity of CG-6-a was even higher and reached 0.662 g/g, which was primarily due to its larger specific surface area and smaller CaO crystal size. After 10 cycles, CG-6-a also maintained a higher CO2 uptake capacity. As shown

Table 1. Crystal Size of Different Samples sample

CG-a

CG-6-a

CC-a

CC-6-a

CA-a

CA-6-a

crystal size (nm)

32.7

15.6

118.3

94.4

124.5

115.3

Figure 3. SEM images of CG calcination at different atmospheres: (a) fresh CG-a, (b) fresh CG-6-a, (c) cycled CG-a, and (d) cycled CG-6-a.

Table 2. Properties of Different Prepared Sorbents sample

surface area (m2/g)

pore volume (cm2/g)

average pore size (nm)

CG-a CG-6-a CA-6-a CC-6-a CG-8-a

14.75 17.50 12.48 13.96 15.86

0.075 0.138 0.053 0.067 0.151

21.5 31.5 15.2 19.2 41.2

(Figure 1a) and DTG (Figure 1b) curves, both CG-air and CGN2 exhibited a major three-step decomposition process. The first loss (stage I) ended at ∼165 °C, and the weight loss was approximately 3.7% when all of the structural water was almost lost. During this stage, both curves almost completely overlapped. The second loss (stage II), which ranged from 165 to 600 °C, was related to the decomposition of CG to intermediate compounds, such as CaCO3, carbon dioxide, carbon monoxide, and acetone (C3H6O), which was also accompanied by the deposition of carbon.39 At this stage, the weight loss under air (71.6%) significantly exceeded that under nitrogen (57.9%),

Figure 4. Cyclic CO2 capture results of CG-a and CG-6-a: (a) evolution of the CO2 capture capacity during 10 cycles and (b) evolution of the carbonation reaction rate at the first cycle. 7545

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Figure 5. Comparison of XRD patterns among different samples: (a) three types of carbon Ca compounds and (b) three types of CaO sorbents.

Figure 6. Comparison of SEM images among different samples: (a) fresh CA-6, (b) fresh CA-6-a, (c) cycled CA-6-a, (d) fresh CC-6, (e) fresh CC-6-a, (f) cycled CC-6-a, (g) fresh CG-6, (h) fresh CG-6-a, and (i) cycled CG-6-a.

in Figure 3c, after 10 cycles, cycled CG-a exhibited severe sintering that prevented CO2 from diffusing into the interior of the sorbent; this, in turn, affects the carbonation reaction. Fortunately, used CG-6-a (Figure 3d) showed only slight sintering, which confirms its better cyclic CO2 capture characteristics. The results of XRD, SEM, N2 adsorption, and cyclic CO2 capture analyses demonstrate that the sorbent calcined in a N2 atmosphere exhibited a smaller CaO crystal size as well as a larger specific surface area and pore volume, which significantly improved its adsorption capacity and stability. 3.2. Effect of the Carbon Source. Figure 1 also compares the pyrolysis diagrams for CA, CC, and CG under a nitrogen atmosphere. Basically, multiple instances of weight loss including three stages (i.e., dehydration, carbon deposition, and carbon CaCO3 compounds decomposition) were observed. The first weight loss of CA under nitrogen (CA-N2) was ∼10%; the second weight loss was ∼32.5%; and the third weight loss of 26.5% occurred at the end at ∼720 °C. The weight loss result is

Figure 7. TEM image of fresh CG-6.

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Figure 8. Cyclic CO2 capture results of CG-6-a, CC-6-a, and CA-6-a: (a) evolution of the CO2 capture capacity during 10 cycles and (b) evolution of the carbonation reaction rate at the first cycle.

Figure 9. Comparison of XRD patterns of different samples: (a) two types of carbon Ca compounds and (b) two types of CaO sorbents.

CC-N2 were 44.40, 0.25, and 20.27%, respectively. Obviously, CG-N2 obtained the highest final organic carbon content. Furthermore, during the carbon deposition stage, the carbonization of CC and CA started later and CG underwent the earliest carbonization, which began at temperatures as low as 200 °C. During the calcium carbonate decomposition stage, CA decomposed the fastest and its decomposition was complete at ∼720 °C. In comparison, CC and CG required higher temperatures to decompose. In particular, CG decomposed the slowest, and its decomposition was not complete until the temperature reached 900 °C. These three organometallic precursors exhibited significantly different pyrolysis processes, which produced diverse microstructured carbon CaCO 3 compounds. Figure 5a compares the phase compositions of the three types of carbon CaCO3 compounds. The main CaCO3 phase was identified in both CC-6 and CA-6. In contrast to CC-6 and CA-6, only a weak amorphous carbon diffraction peak was observed in CG-6. The absence of CaCO3 suggested that CaCO3 in this compound was trapped inside carbon and may have formed nanosized structures. Morishita et al.42 reported that a lower carbonization temperature resulted in the formation of carbon CaCO3 compounds with a smaller crystal size. CG exhibited the earliest carbonization process, which resulted in the production of the smallest crystal size of CaCO3. The SEM images (Figure 6) confirmed the formation of nanostructured carbon CaCO3 compounds in CG-6. As shown in Figure 6a, CA-6 possessed a heterogeneous and relatively coarse structure. CC-6 (Figure 6d) also exhibited an angular morphology and was composed of small fragments. Substantial morphological changes occurred in CG-6

Figure 10. Comparison of SEM images among different samples: (a) fresh CG-8-a, (b) fresh CG-9-a, (c) cycled CG-8-a, and (d) cycled CG9-a.

similar to that reported by Hewitt et al.,40 who also studied the pyrolysis of calcium acetate monohydrate under a nitrogen atmosphere. For CC under nitrogen (CC-N2), the first weight loss was ∼6.1%, the second weight loss was 24.1%, and the third weight loss of ∼31.0% occurred at the end. This result is similar to that reported by Zhou et al.,41 who studied the pyrolysis of calcium citrate tetrahydrate under a nitrogen atmosphere. According to the initial organic content and the final weight (Figure 1a), the final organic contents of CG-N2, CA-N2, and 7547

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Figure 11. Cyclic CO2 capture results of CG-8-a and CG-9-a: (a) evolution of the CO2 capture capacity during 10 cycles and (b) evolution of the carbonation reaction rate at the first cycle.

Figure 8, CG-6-a exhibited the highest initial adsorption capacity and adsorption rate, which resulted from the smallest crystal sizes and porous texture (larger specific surface area and pore volume) of fresh CG-6-a. After 10 cycles, CA-6-a (Figure 6c) and CC-6-a (Figure 6f) both exhibited a substantial sintered morphology. In particular, CA-6-a suffered from the most sintered texture with large and dense crystals. As a result, the most rapid loss in the CO2 uptake capacity was observed for CA-6-a. In comparison, CG-6-a exhibited a higher CO2 adsorption capacity, which indicates its higher durability as a result of its relatively loose morphology, as shown in Figure 6i. Among the three organometallic precursors, CG began to carbonize at the lowest temperature and formed the template with the highest organic carbon content, which resulted in the slowest decomposition of CaCO3 that effectively prevented crystallite agglomeration and created a porous texture. 3.3. Effect of the Carbonization Temperature. CG was further carbonized at higher temperatures (800 and 900 °C) under a nitrogen atmosphere. The carbon Ca compounds (CG-8 and CG-9) and CaO sorbents (CG-8-a and CG-9-a) were structurally characterized. As shown in Figure 9a, no diffraction peaks corresponding to CaO or CaCO3 were observed in CG-8, which is similar to the results observed for CG-6. When the carbonization temperature was increased (900 °C), the CaO phase was identified in CG-9. The different phase compositions resulted in CaO sorbents with different structures. The results shown in Figure 9b revealed that the crystal size of CaO in CG-8a was 28.2 nm, which is close to that in CG-6-a (15.6 nm). Unfortunately, the crystal size of CaO in CG-9-a increased substantially to 126.8 nm. Similar to CG-6-a, the SEM images (Figure 10) confirmed that CG-8-a possessed a blade-like structure (Figure 10a). In contrast, the blade-like architecture disappeared in CG-9-a (Figure 10b). The cyclic CO2 capture characteristics of CG-8-a are shown in Figure 11. Its cyclic performance was not significantly different from that of CG-6-a. Moreover, used CG-8-a did not exhibit a clear sintered morphology (Figure 10c). Although the carbonization temperature increased by 200 °C, the adsorption capacity and stability of the sorbent remained rather high. However, this advantage disappeared at a high temperature (i.e., 900 °C). For CG-9-a, the adsorption capacity of the first cycle was only 0.360 g/g. After 10 cycles, the adsorption capacity decreased rapidly to 0.144 g/g, which was confirmed by the severe sintered morphology of used CG-9-a (Figure 10d).

(Figure 6g), which appeared as a regular blade-like structure (∼20 nm thick). The high-magnification SEM image (inserted picture) and the TEM image (shown in Figure 7) further suggested that ∼50 nm CaCO3 (dark particles) were dispersed over carbon (light particles), which results in a weak amorphous carbon diffraction peak in the XRD patterns. On the basis of the elemental carbon analysis (Table S1 of the Supporting Information), the organic carbon contents of CG-6, CC-6, and CA-6 were 40.19, 21.97, and 0.51%, respectively. Among the three carbon CaCO3 compounds, CG-6 obtained the highest organic carbon content. Therefore, CG-N2 exhibited the slowest decomposition behaviors. This result is supported by the results reported by Blamey et al. and Materic and Smedley43,44 These authors observed that the surface of calcium hydroxide that decomposed in a CO2 atmosphere was covered by a CaCO3 layer that is similar to carbon in this study. This layer suppressed the decomposition of calcium hydroxide. More importantly, slower decomposition could be beneficial for the production of highly porous CaO sorbents. To confirm this hypothesis, we investigated the structures of three types of CaO sorbents by calcination of their respective carbon Ca compounds in air. As shown in panels b and e of Figures 6, both CA-6-a and CC-6-a possessed honeycomb-like morphologies. In contrast to the other samples, CG-6-a (Figure 6h) possessed a blade-like structure, which is remarkably similar to that of CG-6 (Figure 6g). This architecture was based on the template of the carbon Ca compounds used. The XRD patterns of CC-6-a, CA-6-a, and CG-6-a (Figure 5b and Table 1) confirmed that CA-6-a and CC-6-a have relatively large CaO crystal sizes (115.3 and 94.4 nm, respectively). However, the crystal CaO size in CG-6-a was only one-seventh that of CA-6-a. The results from N2 adsorption (Table 2) revealed that the specific surface area, pore volume, and average pore size of CG-6a were higher than those of CA-6-a and CC-6-a. According to the study by Lu et al.,45 during the decomposition of organometallic precursors, a large release of gas facilitated the formation of a porous structure. As shown in Figure S1 of the Supporting Information, during the pyrolysis process (stages I and II), CG exhibited the largest weight loss (∼6.9) that is relative to the initial Ca content, indicating the release of the largest amount of gas. On the other hand, CG-6 obtained the highest organic carbon content. Burn-off of such organic carbon in the secondary calcination step under air provided additional pores. Therefore, CG-6-a may be used to achieve the most porous structure. The cyclic CO2 capture performance of CC-6-a, CA-6-a, and CG-6-a was also investigated. As shown in panels a and b of 7548

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On the important roles of energy integration and sorbent behavior. Appl. Energy 2016, 162, 787−807. (5) Hanak, D. P.; Anthony, E. J.; Manovic, V. A review of developments in pilot-plant testing and modelling of calcium looping process for CO2 capture from power generation systems. Energy Environ. Sci. 2015, 8 (8), 2199−2249. (6) Wang, J.; Huang, L.; Yang, R.; Zhang, Z.; Wu, J.; Gao, Y.; Wang, Q.; O’Hare, D.; Zhong, Z. Recent advances in solid sorbents for CO2 capture and new development trends. Energy Environ. Sci. 2014, 7 (11), 3478−3518. (7) Liu, W. Q.; An, H.; Qin, C. L.; Yin, J. J.; Wang, G. X.; Feng, B.; Xu, M. H. Performance Enhancement of Calcium Oxide Sorbents for Cyclic CO2 Capture-A Review. Energy Fuels 2012, 26 (5), 2751−2767. (8) Manovic, V.; Anthony, E. J. Sintering and Formation of a Nonporous Carbonate Shell at the Surface of CaO-Based Sorbent Particles during CO2-Capture Cycles. Energy Fuels 2010, 24 (10), 5790−5796. (9) Liu, W.; González, B.; Dunstan, M. T.; Saquib Sultan, D.; Pavan, A.; Ling, C. D.; Grey, C. P.; Dennis, J. S. Structural evolution in synthetic, Ca-based sorbents for carbon capture. Chem. Eng. Sci. 2016, 139, 15−26. (10) Ridha, F. N.; Lu, D. Y.; Symonds, R. T.; Champagne, S. Attrition of CaO-based pellets in a 0.1 MWth dual fluidized bed pilot plant for post-combustion CO2 capture. Powder Technol. 2016, 291, 60−65. (11) Barker, R. The reactivity of calcium oxide towards carbon dioxide and its use for energy storage. J. Appl. Chem. Biotechnol. 1974, 24 (4−5), 221−227. (12) Gupta, H.; Fan, L. S. Carbonation-calcination cycle using high reactivity calcium oxide for carbon dioxide separation from flue gas. Ind. Eng. Chem. Res. 2002, 41 (16), 4035−4042. (13) Lan, P.; Wu, S. Mechanism for self-reactivation of nano-CaObased CO2 sorbent in calcium looping. Fuel 2015, 143, 9−15. (14) Zhou, Z.; Qi, Y.; Xie, M.; Cheng, Z.; Yuan, W. Synthesis of CaObased sorbents through incorporation of alumina/aluminate and their CO2 capture performance. Chem. Eng. Sci. 2012, 74 (0), 172−180. (15) Arias, B.; Grasa, G. S.; Alonso, M.; Abanades, J. C. Postcombustion calcium looping process with a highly stable sorbent activity by recarbonation. Energy Environ. Sci. 2012, 5 (6), 7353−7359. (16) Martinez, I.; Grasa, G.; Murillo, R.; Arias, B.; Abanades, J. C. Evaluation of CO2 Carrying Capacity of Reactivated CaO by Hydration. Energy Fuels 2011, 25 (3), 1294−1301. (17) Manovic, V.; Anthony, E. J. Thermal Activation of CaO-Based Sorbent and Self-Reactivation during CO2 Capture Looping Cycles. Environ. Sci. Technol. 2008, 42 (11), 4170−4174. (18) Valverde, J. M.; Sanchez-Jimenez, P. E.; Perez-Maqueda, L. A. Relevant Influence of Limestone Crystallinity on CO2 Capture in The Ca-Looping Technology at Realistic Calcination Conditions. Environ. Sci. Technol. 2014, 48 (16), 9882−9889. (19) Li, Z. S.; Liu, Y.; Cai, N. S. Understanding the effect of inert support on the reactivity stabilization for synthetic calcium based sorbents. Chem. Eng. Sci. 2013, 89, 235−243. (20) Luo, C.; Zheng, Y.; Ding, N.; Wu, Q.; Bian, G.; Zheng, C. Development and Performance of CaO/La2O3 Sorbents during Calcium Looping Cycles for CO2 Capture. Ind. Eng. Chem. Res. 2010, 49 (22), 11778−11784. (21) Peng, W.; Xu, Z.; Luo, C.; Zhao, H. Tailor-Made Core-Shell CaO/TiO2-Al2O3 Architecture as a High-Capacity and Long-Life CO2 Sorbent. Environ. Sci. Technol. 2015, 49 (13), 8237−8245. (22) Tian, S.; Jiang, J.; Yan, F.; Li, K.; Chen, X. Synthesis of Highly Efficient CaO-Based, Self-Stabilizing CO2 Sorbents via StructureReforming of Steel Slag. Environ. Sci. Technol. 2015, 49 (12), 7464− 7472. (23) Wang, K.; Han, D.; Zhao, P.; Hu, X.; Yin, Z.; Wu, D. Role of MgxCa1−xCO3 on the physical−chemical properties and cyclic CO2 capture performance of dolomite by two-step calcination. Thermochim. Acta 2015, 614, 199−206. (24) Wang, K.; Yin, Z.; Zhao, P.; Han, D.; Hu, X.; Zhang, G. Effect of Chemical and Physical Treatments on the Properties of a Dolomite Used in Ca Looping. Energy Fuels 2015, 29 (7), 4428−4435.

4. CONCLUSION The pyrolysis behaviors and resulting carbon Ca compounds played a critical role in the physicochemical properties of the CaO sorbents as well as their cyclic CO2 capture characteristics. In comparison to direct calcination in air, the carbon template route produced a smaller CaO crystal size as well as a larger specific surface area and pore volume. Among the three organometallic calcium compounds, CG began to carbonize at the lowest temperature and formed the template with the highest carbon content, which resulted in the slowest decomposition of CaCO3. These pyrolysis behaviors effectively prevented crystallite agglomeration and sintering and created the most porous texture. Therefore, the most favorable structure was produced, which contributed to the fastest sorption rate, highest adsorption capacity, and best stability of CG-6-a. When the carbonization temperature was between 600 and 800 °C, this outstanding CO2 performance was maintained because the morphology did not change significantly. However, when the carbonization temperature exceeded 900 °C, CaCO3 in the carbon Ca compounds completely decomposed to CaO. Therefore, the carbon template lost its effectiveness, resulting in a rapid loss in the CO2 uptake capacity.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.energyfuels.6b01223. Detailed pyrolysis diagrams (Figure S1), comparison of XRD patterns among CG-a, CC-a, and CA-a (Figure S2), organic carbon contents of several samples (Table S1), and discussion regarding the organic carbon contents (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*Telephone/Fax: +86-516-83592000. E-mail: liyimincumt@ 163.com. *Telephone/Fax: +86-516-83592000. E-mail: wangkecumt@ 163.com. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Fundamental Research Funds for the Central Universities (2015XKMS056). REFERENCES

(1) Jacobson, M. Z. Review of solutions to global warming, air pollution, and energy security. Energy Environ. Sci. 2009, 2 (2), 148− 173. (2) Kierzkowska, A. M.; Pacciani, R.; Müller, C. R. CaO-Based CO2 Sorbents: From Fundamentals to the Development of New, Highly Effective Materials. ChemSusChem 2013, 6 (7), 1130−1148. (3) Boot-Handford, M. E.; Abanades, J. C.; Anthony, E. J.; Blunt, M. J.; Brandani, S.; Mac Dowell, N.; Fernandez, J. R.; Ferrari, M. C.; Gross, R.; Hallett, J. P.; Haszeldine, R. S.; Heptonstall, P.; Lyngfelt, A.; Makuch, Z.; Mangano, E.; Porter, R. T. J.; Pourkashanian, M.; Rochelle, G. T.; Shah, N.; Yao, J. G.; Fennell, P. S. Carbon capture and storage update. Energy Environ. Sci. 2014, 7 (1), 130−189. (4) Perejon, A.; Romeo, L. M.; Lara, Y.; Lisbona, P.; Martinez, A.; Manuel Valverde, J. The Calcium-Looping technology for CO2 capture: 7549

DOI: 10.1021/acs.energyfuels.6b01223 Energy Fuels 2016, 30, 7543−7550

Article

Energy & Fuels (25) Wang, S.; Fan, S.; Fan, L.; Zhao, Y.; Ma, X. Effect of Cerium Oxide Doping on the Performance of CaO-Based Sorbents during Calcium Looping Cycles. Environ. Sci. Technol. 2015, 49 (8), 5021−5027. (26) Zhao, M.; Bilton, M.; Brown, A. P.; Cunliffe, A. M.; Dvininov, E.; Dupont, V.; Comyn, T. P.; Milne, S. J. Durability of CaO-CaZrO3 Sorbents for High-Temperature CO2 Capture Prepared by a Wet Chemical Method. Energy Fuels 2014, 28 (2), 1275−1283. (27) Zhao, M.; Shi, J.; Zhong, X.; Tian, S.; Blamey, J.; Jiang, J.; Fennell, P. S. A novel calcium looping absorbent incorporated with polymorphic spacers for hydrogen production and CO2 capture. Energy Environ. Sci. 2014, 7 (10), 3291−3295. (28) Lysikov, A. I.; Salanov, A. N.; Okunev, A. G. Change of CO2 carrying capacity of CaO in isothermal recarbonation-decomposition cycles. Ind. Eng. Chem. Res. 2007, 46 (13), 4633−4638. (29) Lu, H.; Reddy, E. P.; Smirniotis, P. G. Calcium Oxide Based Sorbents for Capture of Carbon Dioxide at High Temperatures. Ind. Eng. Chem. Res. 2006, 45 (11), 3944−3949. (30) Liu, W.; Feng, B.; Wu, Y.; Wang, G.; Barry, J.; Diniz da Costa, J. C. Synthesis of Sintering-Resistant Sorbents for CO2 Capture. Environ. Sci. Technol. 2010, 44 (8), 3093−3097. (31) Liu, W.; Low, N. W. L.; Feng, B.; Wang, G.; Diniz da Costa, J. C. Calcium Precursors for the Production of CaO Sorbents for Multicycle CO2 Capture. Environ. Sci. Technol. 2010, 44 (2), 841−847. (32) Li, Y. J.; Zhao, C. S.; Chen, H. C.; Liang, C.; Duan, L. B.; Zhou, W. Modified CaO-based sorbent looping cycle for CO2 mitigation. Fuel 2009, 88 (4), 697−704. (33) Ridha, F. N.; Manovic, V.; Wu, Y.; Macchi, A.; Anthony, E. J. Postcombustion CO2 capture by formic acid-modified CaO-based sorbents. Int. J. Greenhouse Gas Control 2013, 16 (0), 21−28. (34) Hu, Y.; Liu, W.; Sun, J.; Li, M.; Yang, X.; Zhang, Y.; Liu, X.; Xu, M. Structurally improved CaO-based sorbent by organic acids for high temperature CO2 capture. Fuel 2016, 167, 17−24. (35) Rodriguez-Navarro, C.; Ruiz-Agudo, E.; Luque, A.; RodriguezNavarro, A. B.; Ortega-Huertas, M. Thermal decomposition of calcite: Mechanisms of formation and textural evolution of CaO nanocrystals. Am. Mineral. 2009, 94 (4), 578−593. (36) Naeem, M. A.; Armutlulu, A.; Broda, M.; Lebedev, D.; Müller, C. R. The development of effective CaO-based CO2 sorbents via a sacrificial templating technique. Faraday Discuss. 2016, DOI: 10.1039/ C6FD00042H. (37) Radfarnia, H. R.; Iliuta, M. C. Limestone Acidification Using Citric Acid Coupled with Two-Step Calcination for Improving the CO2 Sorbent Activity. Ind. Eng. Chem. Res. 2013, 52 (21), 7002−7013. (38) Wang, K.; Hu, X.; Zhao, P.; Yin, Z. Natural dolomite modified with carbon coating for cyclic high-temperature CO2 capture. Appl. Energy 2016, 165, 14−21. (39) Labuschagne, F.; Focke, W. W. Metal catalysed intumescence: characterisation of the thermal decomposition of calcium gluconate monohydrate. J. Mater. Sci. 2003, 38 (6), 1249−1254. (40) Hewitt, F.; Rhebat, D. E.; Witkowski, A.; Hull, T. R. An experimental and numerical model for the release of acetone from decomposing EVA containing aluminium, magnesium or calcium hydroxide fire retardants. Polym. Degrad. Stab. 2016, 127, 65−78. (41) Zhou, Q. Q.; Chen, X. Y.; Wang, B. An activation-free protocol for preparing porous carbon from calcium citrate and the capacitive performance. Microporous Mesoporous Mater. 2012, 158, 155−161. (42) Morishita, T.; Tsumura, T.; Toyoda, M.; Przepiorski, J.; Morawski, A. W.; Konno, H.; Inagaki, M. A review of the control of pore structure in MgO-templated nanoporous carbons. Carbon 2010, 48 (10), 2690−2707. (43) Blamey, J.; Lu, D. Y.; Fennell, P. S.; Anthony, E. J. Reactivation of CaO-Based Sorbents for CO2 Capture: Mechanism for the Carbonation of Ca(OH)2. Ind. Eng. Chem. Res. 2011, 50 (17), 10329−10334. (44) Materic, V.; Smedley, S. I. High Temperature Carbonation of Ca(OH)2. Ind. Eng. Chem. Res. 2011, 50 (10), 5927−5932. (45) Lu, H.; Khan, A.; Smirniotis, P. G. Relationship between Structural Properties and CO2 Capture Performance of CaO-based Sorbents Obtained from Different Organometallic Precursors. Ind. Eng. Chem. Res. 2008, 47 (16), 6216−6220. 7550

DOI: 10.1021/acs.energyfuels.6b01223 Energy Fuels 2016, 30, 7543−7550