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Facile Synthesis of High Surface Area Mesoporous MgO with Excellent High Temperature CO Adsorption Potential 2

Aamir Hanif, Soumen Dasgupta, and Anshu Nanoti Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.6b00647 • Publication Date (Web): 28 Jun 2016 Downloaded from http://pubs.acs.org on June 29, 2016

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Industrial & Engineering Chemistry Research

Facile Synthesis of High Surface Area Mesoporous MgO with Excellent High Temperature CO2 Adsorption Potential Aamir Hanif*a,b, Soumen Dasguptaa,b and Anshu Nanotia,b a

Academy of Scientific and Innovative Research, CSIR-Indian Institute of Petroleum Campus (CSIR-IIP), Dehradun-248005, India. b

Adsorption and Membrane Separation Laboratory, CSIR-Indian Institute of Petroleum, Dehradun-248005, India.

Mesoporous magnesium oxide of high surface area (>350 m2/g) has been synthesized by a simple ammonia precipitation method from aqueous magnesium nitrate solution. This material has been evaluated for CO2 adsorption in the temperature and pressure regime relevant to precombustion capture in integrated gasification combined cycle (IGCC) scheme. CO2 adsorption capacities were also compared with magnesium oxides obtained by other synthetic routes such as hydrothermal urea hydrolysis and direct thermal degradation of magnesium nitrate respectively. The mesoporous magnesium oxide shows CO2 capacity at high temperatures which is comparable or even exceeding the capacity values for this class of high temperature CO2 adsorbents reported in literature so far. The materials studied have been characterized by various characterization techniques like PXRD, surface area, pore size distribution analysis, SEM and

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FT-IR. A cyclic adsorption desorption study was also carried out to understand multi cycle stability of the adsorbent for high temperature CO2 capacity.

Introduction The increasing concentration of carbon dioxide (CO2) in the atmosphere due to burning of fossil fuels is recognized to be a major contributor to the anthropogenic global warming phenomenon.1 The current CO2 level2 in the atmosphere is around 400 ppm much higher than the upper safe limit3 of 350 ppm. Controlling further escalation of CO2 levels is thus a stupendous technological challenge facing mankind. Various strategies such as use of more energy efficient industrial processes and technologies, replacement of the conventional sources of energy with renewable sources and CO2 capture with sequestration have been proposed as a multi-pronged short to long term strategy to tackle this global issue.4 From a more practical perspective however severing our total dependence on fossil fuel based energy seems not to be a feasible option in the short term particularly for developing countries. In this context CO2 capture from the point sources could be a more immediate response.4,5 In order to capture CO2 economically, the point sources producing maximum amount of anthropogenic CO2 needs to be targeted. Electricity and heat production sector is the major source of CO2 production accounting to about 42% and 53% of total anthropogenic CO2 emissions of the World and India respectively.6 Various strategies have been suggested to capture CO2 which can be classified into three main groups viz post combustion capture, pre combustion capture and oxy-fuel combustion.7 The post combustion is end of pipe capture where the gas coming out after combustion in coal based power plants typically contains 10-15% of CO2, 70-75% N2, 8-15% H2O, 3-4% O2 besides traces of SOx, NOx and other gases.8 The temperature of post combustion flue gas varies from 50-150oC and pressure is around 1 atm. In


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pre-combustion capture, which can take place in IGCC cycle, the feed stock (crushed coal) is gasified with steam at very high temperature of 800-900oC. The gasification produces synthesis (syn) gas stream which is a mixture of CO and H2 along with the considerable amounts of CO2, H2S and water vapor. The gasification stream is further subjected to catalytic water gas shift reaction after which the shifted stream may contain 35-50% CO2 along with hydrogen in bulk besides water vapor and H2S. The temperature of shifted water gas stream ranges from 250400oC and pressure about 30-40 bar.9 Whether a pre-combustion or a post combustion CO2 capture route to be followed, the choice of the capture media plays a pivotal role in the overall efficiency of the process. In the context of pre combustion capture the sorbent materials which can capture CO2 at high temperatures are suggested to be materials of choice. These materials should further have a stable capacity over multiple cycles and better kinetics of adsorption besides being economically favourable.10 Various sorbent materials have been tested for high temperature CO2 capture such as calcium oxide,11 lithium silicates,12,13 lithium zirconates,14 magnesium oxide15 and magnesium aluminum hydrotalcites.16–18 Calcium oxide based sorbents have very high CO2 capacity exceeding 11 mmol/g at 1 bar CO2 pressure and temperature of 600-700oC. Their regeneration temperature however, is very high (900-1000oC) implying requirement of large temperature swing between sorption and regeneration. Also the kinetics of CO2 adsorption is found to be poor in the operational temperature window of IGCC i.e. 200-500oC.19 Further rapid decay of the CO2 capacities was also observed over multiple cycles due to gradual sintering of the adsorbent.11 Similarly lithium silicates and zirconates, despite having moderately high CO2 capacities up to 6.5 mmol/g at 1 bar pressure and 500-600oC temperature,19 have also slower CO2 uptake kinetics20 in the operating temperature regime of IGCC in addition to the requirement of large



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temperature swing for regeneration. Hydrotalcite derived magnesium aluminum mixed oxides although having a disadvantage of lower CO2 capacity are still receiving attention since these materials have better CO2 adsorption kinetics,21 can be regenerated with lower energy penalty and retain cyclic capacity over multiple cycles of adsorption and regeneration.21,22 Our previous work in this domain focuses on improvement of the capacity of magnesium aluminum mixed oxides by following various modified synthesis strategies.22,23 One of the intriguing facts about magnesium aluminum mixed oxides is that the aluminum oxide phase has minimum interaction with CO2 and the active phase responsible for the adsorption of CO2 is mainly magnesium oxide.24 Aluminum oxide phase is said to enhance the capacity by maintaining the small magnesium oxide particles in dispersed state and thus removing the diffusion limitations associated with bulk magnesium oxide.24,25 This suggests that high surface area nano-sized magnesium oxide can also be a promising candidate for high temperature CO2 capture. Various attempts have been made for synthesizing high surface area magnesium oxides such as aerogel synthesis,26 chemical vapor deposition,27 template synthesis,28 and thermal degradation29–31 methods. Aerogel and chemical vapor deposition produce ultrahigh surface area magnesium oxide (350-500 m2/g BET) but require drastic conditions of temperature and pressure during synthesis besides organic solvents and expensive instrument.26,27 Hard template synthesis afford moderately high surface area (up to 300 m2/g BET) mesoporous magnesium oxides but require multiple reagents and synthesis steps making the whole process less attractive.28 More recently thermal degradation of magnesium compounds has shown promise to be a simple method to produce high surface area mesoporous magnesium oxides (250-300 m2/g). However successful results were obtained only with economically less attractive organo magnesium compounds32 such as magnesium acetate as precursors and that too without completely eliminating the need of


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other organic solvents33 or additionally requiring inert purge during the thermal degradation process.31 Keeping all these drawbacks of different synthesis procedures in mind we report here a facile synthesis of high surface area (>350 m2/g) nano-sized mesoporous magnesium oxide by ammonium hydroxide assisted precipitation of magnesium hydroxide from magnesium nitrate aqueous solution followed by thermal degradation of the precipitated magnesium hydroxide in vacuum. The magnesium oxide thus synthesized has been evaluated for CO2 capture in the temperature range of 300-400oC and pressures up to 19 bar since these pressure and temperature conditions are typically encountered in the IGCC scheme. The synthesized material has also been compared for CO2 uptake capacity and kinetics with the magnesium oxide samples obtained from magnesium nitrate by the urea hydrolysis and thermal degradation method respectively. The optimized adsorbent has also been compared with the CO2 capacities of other reported magnesium oxides and hydrotalcites at relatively lower pressures and temperature since no data is available in literature at higher pressure (>1 bar) for these type of adsorbents Experimental Synthesis The synthesis of magnesium oxide from magnesium nitrate precursor was done by three methods namely ammonia precipitation, urea hydrolysis and thermal degradation. Ammonia precipitation: A 4M aqueous solution of magnesium nitrate hexahydrate was taken in a glass bottle at room temperature. A calculated amount of aqueous ammonium hydroxide solution, so as to maintain NH4OH to Mg mole ratio of 5, was added to the magnesium nitrate solution and the mixture was stirred magnetically for about 5 minutes at room temperature. The bottle was then tightly closed and gradually heated to 60oC and maintained at this temperature



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for 6 hours with continuous stirring. After six hours the cap of the reagent bottle was slowly opened. A brisk effervescence was observed due to release of accumulated ammonia gas inside the bottle. The heating was stopped and slow stirring continued for next 24 hours. The precipitate was then separated by centrifugation and repeatedly washed with distilled water several times. The precipitate was dried at 70oC for another 24 hours. The solid obtained was coded as MGOA. A part of this material was calcined at 400oC under vacuum for 12 hours. The calcination temperature was attained at a very slow rate of 1oC/min. The calcined sample was coded as MGO-AC. Urea hydrolysis: Here calculated amount of urea was added to 4M magnesium nitrate hexahydrate aqueous solution, so as to maintain urea to Mg mole ratio of 5, in a teflon lined autoclave with vigorous stirring so as to dissolve urea completely. The autoclave was next closed and placed in a preheated oven at 90oC for six hours. The heating was stopped after 6 hours and the sample was removed from autoclave after 24 hours. Further processing and calcination of the sample was done following procedure similar to the ammonia precipitated sample. The as synthesized sample and its calcined form were coded as MGO-U and MGO-UC respectively. Thermal degradation method: In this synthesis method magnesium nitrate hexahydrate salt was directly calcined at 400oC under vacuum for 12 hours. Here also a gradual temperature rise of 1oC/ min was applied for achieving the final calcination temperature. The solid thus obtained was coded as MGO-TC. Characterization The synthesized samples were characterized by PXRD, FT-IR, surface area, pore size distribution analysis and SEM techniques. PXRD scans were carried out in Bruker (Germany) advanced diffractometer using Cu Kα radiation in the 2θ range of 5-80o at a scan rate of 2o/min.


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FT-IR (Nicolet 8700) scans of CO2 saturated magnesium oxide were taken after adsorbing CO2 at 300oC temperature and 1 bar pressure in the adsorption cell of Hiden IGA-001 gravimetric microbalance unit. About 1mg of the material was mixed with 100mg of KBr from which pellet was made for FT-IR analysis. Surface area and pore size distribution analysis was carried out in Micromeritics Tristar surface area and porosity analyzer. Prior to surface area determination the samples were degassed at 400oC under helium flow for 10 hours. SEM micrographs were taken by Quanta (Netherlands) 200F FESEM. Before SEM analysis the adsorbents were vacuum dried at 70oC. Equilibrium isotherm, CO2 uptake rate and cyclic adsorption capacity measurements These measurements were carried out in a gravimetric microbalance unit (Hiden IGA-001, UK). About 50 mg of adsorbent was loaded in the sample holder of the instrument and activated at 400oC under evacuation till their weight became constant. After this the sample holder was cooled under evacuation to the target temperature of isotherm measurement, maintained by an external thermostat. Ultrapure CO2 (99.999 mole%, Chemtron, Navi Mumbai) was then introduced at various pressures ranging from near vacuum to 19 bar to get the isotherm in this pressure range. At each pressure point enough time was given to achieve equilibrium. In this manner the adsorption isotherms up to 19 bar were obtained for different adsorbents. The kinetic CO2 uptake measurements were carried out at 300oC and 60 millibar pressure. The low pressure was chosen to minimize time lag in attaining the final pressure during pressure dosing. For monitoring the kinetics of adsorption with greater accuracy at a particular pressure the pressure should be attained within minimum time so that the gas uptake during pressure change is very small and can be neglected. In our case the 60 millibar pressure was attained in less than 7 seconds which is negligible compared to the 1-hour duration during which the CO2 uptake



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occurred for 1 hour till further weight uptake became negligible. For cyclic studies after every adsorption step the adsorbent was regenerated at 400oC under high vacuum (10-6 millibar) till it regained its original and stable weight as before adsorption. The adsorbent was then cooled back to its adsorption temperature under vacuum before another cycle of adsorption took place. Results and discussion The synthesis of magnesium oxide by direct thermal decomposition of magnesium nitrate proceeds with evolution of water, NO2 and O2.34 PXRD scans (Figure 1) of the calcined and noncalcined samples reveal that during urea hydrolysis method magnesium carbonate hydroxide hydrate [MGO-U: Mg5(CO3)4(OH)2.(H2O)4, JCPDS:70-1177]

is formed which then gets

converted to magnesium oxide [MGO-UC: JCPDS:87-0651] during calcination whereas in ammonia precipitation magnesium hydroxide [MGO-A: JCPDS:84-2163] is formed as an intermediate which on calcination forms magnesium oxide [MGO-AC: JCPDS:87-0651] by evolution of water vapor. However, the crystallinity of magnesium oxide varies greatly depending on synthesis methods followed. The intensity of the diffraction peaks decreases in the order MGO-TC>MGO-UC>MGO-AC which indicates that magnesium oxide formed by direct thermal degradation (MGO-TC) is highly crystalline compared to those synthesized through the urea hydrolysis (MGO-UC) and ammonia precipitation route (MGO-AC). PXRD data is further substantiated by SEM analysis which shows well-developed octahedral crystallites having diagonal dimensions of above 1.2 micron in the sample MGO-TC. In case of MGO-UC a sheet like morphology is visible whereas in MGO-AC irregular disrupted platelet type morphology is seen (Figure 2). The sheet like morphology of the magnesium oxide sample MGO-UC is similar to that of its precursor magnesium carbonate hydroxide hydrate (MGO-U) which shows that the morphology


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is almost intact after calcination. It may be mentioned here that the magnesium carbonate hydroxide hydrate is known to adopt various such morphology under hydrothermal conditions depending upon temperature and pH.35 With regard to the magnesium oxide sample MGO-AC a disrupted platelet morphology can be seen. This morphology is reminiscent of the brucite structure of Mg(OH)2 formed as intermediate during MgO synthesis by ammonia precipitation method as evident from the SEM micrographs of these samples. Prior to calcination agglomerated disc shaped crystallites of Mg(OH)2 is visible (MGO-A). Such morphology may be the result of the weak hydrogen bonding between layers of magnesium hydroxide36 which during calcination gets converted to MgO with the breakage of hydrogen bonding with subsequent release of water vapors leading to disrupted irregular platelet type morphology. Moreover, a more controlled growth of the magnesium carbonate hydroxide hydrate structures as the pH level of the mother liquor rises gradually during the slow hydrothermal decomposition of urea may have led to larger sheet like morphology. In contrast a much faster rate of nucleation of magnesium hydroxide during the ammonia assisted precipitation is probably responsible for the formation of smaller nano crystallites of the brucite structure as observed in MGO-A. Similar observation was also made during precipitation synthesis of structurally related Mg-Al hydrotalcites.37,38 FT-IR spectra of MGO-UC and MGO-AC (Figure 3) show an intense peak at 3700cm-1 and abroad peak centered around 3430cm-1 indicative of the presence of residual surface hydroxyl (OH) groups39 even after calcination of the hydroxide from MGO-U and MGO-A respectively. No such well differentiated peaks are observed in MGO-TC since that is formed by the direct calcination of magnesium nitrate and thus no residual -OH group is present. A negligible



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intensity hump in the 3400-3700cm-1 region may be due to presence of small amount of surface moisture in the sample. The nitrogen adsorption isotherms at liquid nitrogen temperature (77K) of the synthesized adsorbents are shown in Figure 4. MGO-AC shows highest adsorption followed by the sample MGO-UC and MGO-TC shows negligible adsorption indicating the difference in surface area of samples. The shape of isotherms of samples MGO-AC and MGO-UC suggests an overall type H3 hysteresis loop indicative of the presence of slit shaped mesopores.40,41

However, a close

observation of the isotherms shows some deviation from the ideal H3 shape and the isotherm actually comprises of two separate segments of hysteresis. Pore size distribution calculated from desorption branch of isotherm42 using BJH model is shown in Figure 5. The material MGO-AC shows a bimodal pore size distribution with a narrow distribution centered around 3nm and a very broad distribution centered around 25nm. This can be as a consequence of the slits of various dimensions and shapes ranging from parallel to wedge shaped slits formed in the disrupted platelet morphology of MGO-AC.42 MGO-UC has smaller pore volume but also shows similar bimodal pore size distributions (Figure 5 inset) centered around 3nm and 18nm whereas the sample MGO-TC has negligible porosity. The surface area and pore volume calculated from these isotherms of the synthesized samples is given in Table 1. The surface area and pore volume of the samples varies as MGO-AC>MGO-UC>MGO-TC. This trend is expected as the sample MGO-AC has smallest crystallite size followed by the MGO-UC and MGO-TC respectively. The smaller crystallites present a larger surface area in comparison to large crystallites. The presence of high degree of porosity in MGO-AC and MGOUC samples could be a result of the evolution of water vapor and adsorbed basic ammonia


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molecules from the magnesium hydroxide surface during the calcination step. In contrast the evolution of acidic NO2 along with water vapor during the thermal degradation of the magnesium nitrate hexahydrate may be detrimental for the creation of pores in magnesium oxide. Thus the magnesium oxide prepared from the thermal degradation of nitrate salt has least pore volume and surface area. In order to assess potential of these materials for high temperature CO2 adsorption, isotherms were measured at 300oC for all these adsorbents up to 19 bar, the conditions of the pressure and temperature typically encountered in pre-combustion CO2 capture catalytic water gas shift reaction during IGCC scheme. These isotherms are shown in Figure 6. The CO2 capacity of these adsorbents follows the order MGO-AC>MGO-UC>MGO-TC. In fact, MGO-TC has much lower capacity in comparison with the other two adsorbents which could be attributed to negligible surface area of this material. As adsorption is primarily a surface phenomenon the smaller crystallites having higher surface area have higher CO2 adsorbing active site density as compared to larger crystallites. The nature of interactions between adsorbent MgO and adsorbate CO2 have been studied by several authors29,43–46 using FT-IR where different peaks in the wave number range 800-1650cm-1 have been ascribed to different modes of adsorbed CO2 on the MgO surface. In order to check the CO2 interactions with our adsorbent samples, FT-IR analysis was carried out at ambient conditions after saturating the adsorbents with CO2 at 300oC. It is pertinent to mention that we could not perform FT-IR analysis at 300oC with our setup. However, FT-IR analysis at ambient temperature after saturating the sample with CO2 at 300oC can be indicative of CO2 interactions with the adsorbent at 300oC, as the gas once adsorbed at high temperature (300oC) will not desorb on cooling to ambient conditions. In order to establish that CO2 remains chemically adsorbed and not desorbed under ambient conditions, a desorption isotherm was



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measured at 25oC after completion of the adsorption isotherm at 300oC on sample MGO-AC in the Hiden IGA-001 gravimetric analyzer. The desorption measurement was carried by first cooling the adsorbent to 25oC from 300oC under same regulated CO2 pressure (19 bar). Next the pressure was decreased in a stepwise manner at the controlled rate of 200 millibar/min from 19 bar to high vacuum (10-6 millibar). In order to achieve equilibrium enough time was given at each desorption pressure point. As evident in Figure 7 CO2 once adsorbed at 300oC didn’t get desorbed at 25oC even after applying high vacuum rather further CO2 uptake took place during cooing. FT-IR analysis of our samples (Figure 3) in CO2 interaction region (800-1650cm-1) clearly indicates that there is least interaction of CO2 with the sample MGO-TC as it has negligible CO2 capacity among the three adsorbent samples. In case of MGO-TC a very small peak at 1390cm-1 was observed which can be assigned to unidentate mode of interaction of carbonate species.43 The samples MGO-UC and MGO-AC however show similar peaks of considerable intensity in this region (800-1650cm-1) due to higher adsorption of CO2 on these adsorbents. A closer look at FT-IR spectra of MGO-AC (Figure 8) shows peaks centered around 862cm-1and 1085cm-1. In addition to these peaks a very broad peak around 1350-1650cm-1 was also observed which was deconvoluted into three peaks centered around 1408cm-1, 1493cm-1and 1629cm-1 respectively. The peak around 1640 cm-1 can be assigned to bidentate mode of interaction and all other peaks are assigned to unidentate mode of CO2 interaction with the magnesium oxide adsorbent on the basis of peak assignments made in earlier reports.46 It has also been reported that the corner and edge sites of the MgO crystallites are most active in CO2 adsorption and the unidentate mode of interaction occurs on the edge sites whereas bidentate mode of interaction occurs on corner sites of these crystallites.47


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Since there are two different types of sites for the CO2 interaction with MgO, the dual site Langmuir (DSL) model48 was used to predict the isotherms. The equation of DSL model can be written as ݊∗ =

݉ଵ ܾଵ ܲ ݉ଶ ܾଶ ܲ + 1 + ܾଵ ܲ 1 + ܾଶ ܲ

where n* is predicted isotherm capacity, P is pressure, m1 and m2 are the monolayer capacities of the two types of sites and b1 and b2 are their respective affinities for CO2. The model fits well with the experimental data as shown in Figure 6 thus further validating the presence of two different types of sites on MgO for CO2 adsorption. The parameters of this model are given in Table 2. CO2 uptake kinetics of the samples MGO-AC and MGO-UC reported in Figure 9 indicates that the average uptake rate of the sample MGO-AC is more than that of sample MGO-UC. A modified equation of double exponential kinetic model reported elsewhere49 was used to predict the kinetic data. The model takes into consideration both types of interactions as observed in our adsorbent. The mathematical equation for this model is: ‫ݍ‬௧ = ‫ݍ‬ଵ ሺ1 − ݁ ି௞భ ௧ ሻ + ‫ݍ‬ଶ ሺ1 − ݁ ି௞మ ௧ ሻ Where q1, q2 are equilibrium capacities at two different types of sites and k1 and k2 their respective rate parameters. The kinetic rate parameters obtained from model fit are given in Table 3. The values of rate parameters of adsorbent MGO-AC are marginally higher than the sample MGO-UC suggesting faster kinetics of adsorption on MGO-AC. In order to study the effect of temperature on adsorption the best performing adsorbent MGO-AC was further tested for the CO2 adsorption at 350oC and 400oC (Figure 10) up to 19 bar as these temperatures are typically encountered in IGCC scheme after catalytic water gas shift reaction. It is observed that the equilibrium capacity of the adsorbent is temperature dependent and



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decreases with the increasing temperature. This behavior is because of decreased affinity of CO2 towards MgO at higher temperatures which is also reflected in the decrease of affinity parameters in DSL model fit of the isotherms. It is worth mentioning that there are few reports of CO2 adsorption on MgO at temperatures exceeding beyond 200oC. However, the related compounds such as Mg-Al hydrotalcites, hydrotalcite derived layered double oxides (LDO), MgAl-O aerogels have been reported for CO2 adsorption at relatively higher temperatures. Further CO2 adsorption data at higher pressures, as required for pre combustion capture application, is scarcely reported with these type of adsorbents. A comparison of CO2 adsorption capacity of our adsorbent MGO-AC with similar type of literature adsorbents under same conditions of temperature and pressure is given in Table 4. It can be observed that the adsorbent MGO-AC reported here has one of the best reported capacities among magnesium oxide based adsorbents for high temperature CO2 adsorption. The cyclic capacity of the optimized adsorbent MGO-AC at 300oC and 1 bar pressure over 27 cycles is shown in Figure 11. It can be observed that there is a rapid fall in capacity by about 38% from 1.71 mmol/g initial capacity to nearly 1.06 mmol/g in first 20 cycles after which the further drop in capacity is very small (4%) up to 27 cycles studied. The initial rapid fall in CO2 capacity of MGO-AC can be ascribed to sintering of the smaller crystallites of the adsorbent. However, over multiple cycles the material finally assumes a stable structure which is somewhat resistant to sintering and hence the rate of fall in further cycles is minimised. This observation is also corroborated by SEM micrographs (Figure 12) of the adsorbent taken before CO2 adsorption and after 27 continuous cycles of CO2 adsorption and regeneration. The freshly calcined adsorbent prior to CO2 adsorption shows a morphology consisting of disrupted platelets along with fluffy structures. Following multiple cycles of CO2 adsorption and desorption the


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fluffy structure of the crystallites disappeared however the disrupted platelet structure remains intact. Conclusions Magnesium oxide was synthesized from magnesium nitrate hexahydrate precursor by three different synthesis strategies namely thermal degradation, urea hydrolysis and ammonia precipitation method. These materials were then compared for high temperature (200-400oC) CO2 adsorption under equilibrium conditions. The synthesis method has bearing on the surface area which in turn affects the CO2 adsorption capacity of the material. Higher surface area improves upon the CO2 capacity of the MgO adsorbent. Unlike the literature reports for synthesis of high surface area nanoporous MgO wherein drastic multistep reaction conditions or costly precursors are used, a facile ammonia precipitation method is proposed in the present study. Further use of aqueous medium and relatively cheaper reagents makes it more beneficial environmentally and economically. The magnesium oxide nanoporous powder thus synthesized has higher CO2 adsorption capacity at high temperatures than most of the reported magnesium based adsorbents. High CO2 capacity at high temperature indicates a strong chemical interactions of CO2 molecules with MgO surface which is also confirmed by desorption as well as by FT-IR studies. The multicycle adsorption desorption experiments indicate that there is an initial drop in the CO2 capacity which then stabilizes gradually. This phenomenon could be due to stabilization of the morphological structure of the adsorbent after initial sintering as also evident from the SEM study. The optimized material thus reported can be a low cost potential sorbent for various high temperature CO2 capture applications such as in IGCC and sorption enhanced water gas shift reaction. Corresponding Author



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*Email: [email protected] , [email protected] Tel: +91-135-2525783 Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Acknowledgements The authors thankfully acknowledge CSIR India for the laboratory facilities. Aamir Hanif is also thankful to UGC- India for UGC-NET research fellowship. Mr. Sandeep Saran, Mr. Raghubir Singh and Mr. L N Sivkumar Konthala are acknowledged for XRD, FT-IR and SEM analysis respectively. References (1)

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Hansen, J.; Sato, M.; Kharecha, P.; Beerling, D.; Berner, R.; Masson-Delmotte, V.; Pagani, M.; Raymo, M.; Royer, D. L.; Zachos, J. C. Target Atmospheric CO2: Where Should Humanity Aim? Open Atmos. Sci. J. 2008, 2, 217.

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Table 1. Surface area and pore volume of the synthesized adsorbents Adsorbent

BET S.A (m2/g)

Pore Volume (cm3/g)

MGO-AC

361.54

0.701

MGO-UC

190.69

0.231

MGO-TC

4.59

0.032



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Table 2. Dual site langmuir model parameters for isotherms on different adsorbents Temperature (oC)

‘b1’ (bar-1)

‘b2’ (bar-1)

200

86.23

0.08

300

33.76

0.05

350

9.43

0.02

400

2.25

0.01

MGO-UC

300

40.66

MGO-TC

300

3.71

Adsorbent

MGO-AC

‘m1’ (mmol/g)

‘m2’ (mmol/g)

1.22

6.13

0.11

0.86

5.35

0.06

0.07

0.38


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Table 3. Double exponential model parameters for kinetic data fitting. Adsorbent

q1(mmol/g)

q2 (mmol/g)

k1 (min-1)

k2 (min-1)

MGO-AC

0.632

0.218

1.606

0.070

MGO-UC

0.521

0.140

1.530

0.036



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Table 4. Comparison of CO2 capacity of adsorbent MGO-AC with other literature reports Adsorbent

Precursors

Synthesis procedure

P (bar)

T (oC)

Capacity (mmol/g)

Reference

MgO

Magnesium Acetate

Carbon insertion/thermal decomposition

1

200

0.66

31

MgO-γAl2O3

Magnesium Nitrate, Aluminum Nitrate

Hydrothermal using P123 micelles

1

200

1.75

50

MgO-Al2O3

Magnesium Nitrate, Aluminum Nitrate

Aerogel procedure

1

200

0.59

51

200

2.00

MgO

Magnesium Nitrate

Precipitation followed by thermal degradation

1

This work 300

1.71

MgO (mesoporous carbon stabilized)

Magnesium Chloride

Fast Pyrolysis method with carbon source

1

300

1.3

52

Mg-Al Hydrotalcite

-

Commercial

1

300

0.52

53


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Figure 1. PXRD of (a) MGO-TC (b) MGO-UC (c) MGO-AC (d) JCPDS: 87-0651 [MgO] (e) MGO-U (f) JCPDS:70-1177[Mg5(CO3)4(OH)2(H2O)4] (g) MGO-A (h) JCPDS:84-2163 [Mg(OH)2]



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Figure 2. SEM micrographs of synthesized samples


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Figure 3. FT-IR of (a) MGO-TC (b) MGO-UC and (c) MGO-AC after CO2 adsorption at 300 oC.



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Figure 4. N2 adsorption isotherms at 77K.


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Figure 5. BJH desorption pore size distribution of synthesized samples. BJH desorption of MGO-UC is also shown inset for clarity.



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Figure 6. CO2 adsorption isotherms of synthesized adsorbents at 300oC. Points denote experimental data and lines dual site langmuir (DSL) model Fit.


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Figure 7. CO2 adsorption and desorption isotherms of MGO-AC.



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Figure 8. FT-IR of adsorbent MGO-AC after adsorption of CO2 at 300oC. The solid line shows experimental data and dotted lines are deconvoluted peaks.


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Figure 9. Experimental CO2 uptake of adsorbents and double exponential model fitting at 60 millibar pCO2 and 300oC temperature.



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Figure 10. Isotherms on adsorbent MGO-AC at various temperatures. Points denote the experimental data and lines dual site langmuir (DSL) model fit isotherms


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Figure 11. Capacity of adsorbent MGO-AC at 300oC and 1 bar CO2 pressure over multiple cycles of adsorption and regeneration.



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Figure 12. SEM micrographs of MGO-AC (a) After calcination but before CO2 adsorption (b) after 27 cycles of adsorption and regeneration.


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