Introduction

Fossil fuel accounts for the world’s major energy supply and its use is anticipated to be continued throughout the 21st century [1]. The use of fossil fuel is always accompanied with a vast emission of CO2. The anthropogenic CO2 emissions upset the natural carbon cycle leading to an increased atmospheric CO2 concentration. No one can deny that there is an urgent need to develop methods for CO2 mitigation.

Currently, a large focus is devoted on capturing CO2 from coal-fired power plant flue streams, which continues to be the major CO2 contributor [2].There are three main technological approaches for CO2 capture- post-, oxy- and pre-combustion. Post-combustion systems capture CO2 from N2-rich flue gas stream, produced by burning fossil fuel in air. Oxy-combustion uses pure stream of O2, instead of air, to combust coal and thus produces CO2-rich gas stream. On the other hand, pre-combustion systems are designed mainly to remove CO2 from the syngas (CO + H2) prior to its combustion for power production [3]. Under pre-combustion conditions, after the water–gas shift reactor \(({\text{CO}} + \;{\text{H}}_{ 2} {\text{O}} \to \;{\text{CO}}_{ 2} \; + \;{\text{H}}_{ 2} )\), the gas stream mainly consists of CO2, H2O and H2. The partial CO2 pressure for pre-combustion capture conditions is around 20–30 bar and the temperature is between 250 and 450 °C [4].

Both the physical and chemical solvents can be used for pre-combustion CO2 capture. Unlike chemical solvent, physical solvent (such as Selexol and Rectisol) selectively absorbs CO2 without forming any chemical bonds. Thus, the physical solvent requires relatively less regeneration energy as compared to that of chemical solvent [5]. However, these physical solvent-based processes suffer severe disadvantages: (1) lose pressure during regeneration step, and (2) require a low operating temperature. Therefore, syngas needs to be cooled prior to CO2 absorption step to attain a relatively low operating temperature. After CO2 absorption, the remaining hydrogen gas stream requires to be reheated to the gas turbine inlet temperature. However, chemical solvents have the advantage of high mass transfer driving force into solution and better acid gas selectivity. Also, chemical solvents can be used in processes that utilize thermal swing regeneration and generate the CO2 at elevated pressure [6]. But chemical solvents increase the energy and cost penalty and thus are downgraded as a future CO2 sorbent [7].

Currently, the focus is to develop advanced physical and chemical solvent systems that have the potential to provide significant improvements in both cost and performance as compared to the Selexol and Rectisol for pre-combustion CO2 capture. The challenges are to modify regeneration conditions to recover the CO2 at a higher pressure, improve selectivity to reduce H2 losses, and develop a solvent that has a high CO2 loading at a higher temperature, which would increase integrated gasification combined cycle (IGCC) efficiency. In the same line, another vital task is to develop a new sorbent which could be highly efficient for pre-combustion capture conditions. Consequently, the US Department of Energy (DOE) performed the thermodynamic modeling activities that included screening analyses for a number of different metal oxides, zirconates, silicates and titanates under various operating conditions to identify new solid sorbents for pre-combustion CO2 capture. About 18 sorbents were modeled and finally seven candidates are chosen—magnesium oxide (MgO), calcium oxide (CaO), lithium zirconate (Li2ZrO3), calcium zirconate (CaZrO3), barium zirconate (BaZrO3), barium titanate (BaTiO3) and barium silicate (BaSiO3) for further investigations [8]. Among these solid sorbents, MgO and CaO are attractive because of their easy accessibility and favorable thermodynamic properties [9].

The alkaline earth metal oxides (such as CaO and MgO) combine with CO2 to form thermodynamically stable carbonates. Metal carbonates, when heated, liberate pure stream of CO2 gas and regenerate the oxides. Eventually, the generated pure CO2 gas can either be sequestered underground or used for enhanced oil recovery [10]. CaO are abundant and thus relatively easily accessible than MgO. However, CaO as CO2 sorbent suffers severe major drawbacks and a detailed discussion can be found elsewhere [11]. It is a fact that regeneration of oxides needs a lot of energy [12]. MgO and Mg(OH)2 are known to be better candidates than CaO for CO2 capture applications due to their low regeneration energy requirement and low operating temperature [13]. Thus, MgO (periclase), although occurs only rarely as an oxide, we choose it here for our study. One should note that MgO when recycled between naturally occurring magnesite or dolomite can cause relatively lesser energy or carbon emission penalty [14].

The CO2 absorption capacity of MgO was studied as a function of particle size, surface area, temperature, pressure, support and concentration of water vapor. The CO2 uptake capacity on different MgO sorbents at different conditions is listed in Table 1. Most of the experiments are restricted to low temperature and ambient or low pressure condition. Hence, the reported sorption capacities of CO2 on MgO are not very high. However, the sorbents which can operate in the range of 300–350 °C would be ideal for the use in IGCC applications. Therefore, we study the CO2 capture capacity of MgO in the relatively high temperature and pressure condition of 300–375 °C and 10–50 bars, respectively.

Table 1 CO2 uptake capacity of MgO obtained from the literature
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Experimental procedure

The reaction

$${\text{MgO}}\left( {\text{s}} \right) \, + {\text{ CO}}_{ 2} \left( {\text{g}} \right) \, \leftrightarrow {\text{ MgCO}}_{ 3} \left( {\text{s}} \right)$$

was studied experimentally. The carbonation reaction was performed in a closed system which permits us to efficiently maintain high temperature and pressure for a long period of time. The decomposition reaction of carbonate was examined using Thermo gravimetric Analysis (TGA) technique.

About 0.2 g of 325-mesh-sized magnesium oxide (delivered by Alfa Aesar) was put inside a closed cylindrical vessel (1.25″ long and 0.35″ internal diameter). About 0.1 ml of water was also introduced on the vessel walls. Chemically pure CO2 gas (Airgas) was passed into this system. Before experiments, CO2 gas was flushed three times to ensure a pure CO2 atmosphere inside the reactor. The reaction was performed for 30 min at a desired temperature and pressure condition. Once the reaction completed, the system was air-cooled. The product was then ground using mortar and pestle. The powder particle was again put back for another reaction at the same experimental condition. This cycle was repeated until no increment in weight of the product was observed.

Thermo gravimetric analysis (TGA) of the product was done using TGA 2950 Thermo gravimetric analyzer. The samples (10 mg) were heated under argon purge, at a heating rate of 10 °C/min to a final temperature of 800 °C. The TGA provides continuous measurements of the sample weight as a function of time and temperature. The amount of formation of MgCO3 was analyzed by the percent loss in weight of the sample while heating up to 800 °C in an argon atmosphere.

The product characterization was performed using X-ray powder diffraction method. Bruker GADDS/D8 is equipped with Apex Smart CCD Detector and direct-drive rotating anode. The MacSci rotating anode (Molybdenum) operates with a 50 kV generator and 20 mA current. X-ray beam size can vary from 50 to 300 µm. The usual collection time is 1200 s.

An isothermal gas adsorption was employed to measure internal surface areas of the powder particles. Micromeritics Tristar II 3020 (surface area and porosimetry analyzer instrument) was used with N2 as adsorptive gas at 77 K (liquid nitrogen bath). The samples were first degassed under 300 °C with a N2 gas flow for 1 h to remove the moisture and other adsorbed gases before analysis. The internal surface area was calculated using the Brunauer–Emmett–Teller (BET) method. The pore volume was also calculated from the adsorbed nitrogen after complete pore condensation (P/P 0 = 0.9925) using the ratio of the densities of liquid and gaseous nitrogen. The pore size was calculated using the Barrett–Joyner–Halenda (BJH) method.

Results and discussion

Figure 1 confirms the formation of MgCO3 at different temperatures and CO2 pressures. MgO, Mg(OH)2, MgO·2MgCO3 and MgCO3 were identified conventionally by their corresponding Joint Committee Powder Diffraction Standard (JCPDS) card number 79-0612,82-2345,31-0804 and 86-2345, respectively.

Fig. 1
figure 1

XRD patterns for MgCO3 formation after reaction at various temperatures and CO2 pressure of a 10 bars and b 50 bars

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One of the vital factors in gas–solid carbonation reaction is the presence of water and there have been numerous observations where water acts as a catalyst [31–34]. Therefore, we have also used water (0.1 ml) for the MgO–CO2 reaction. In absence of water, no CO2 was absorbed at these conditions due to the kinetic limitations. The CO2 sorption capacity of MgO increases significantly in the presence of water vapor. Under humid condition, MgO rapidly locks CO2 in the form of MgCO3. Recently, Fagerlund et al. [29] proposed the reaction mechanism for MgO carbonation in the presence of steam:

$${\text{MgO }} + {\text{ H}}_{ 2} {\text{O }} \leftrightarrow {\text{ MgO}} \cdot {\text{H}}_{ 2} {\text{O}}^{*}$$
$${\text{MgO}} \cdot {\text{H}}_{ 2} {\text{O}}^{*} \, + {\text{ CO}}_{ 2} \leftrightarrow {\text{ MgCO}}_{ 3} + {\text{ H}}_{ 2} {\text{O}}$$
$${\text{MgO }} + {\text{ CO}}_{ 2} \leftrightarrow {\text{ MgCO}}_{ 3}$$

Figure 2 depicts the adsorption/absorption model for MgO–CO2–H2O reaction. Based on the previous work, it could be reasonable to corroborate that water vapor surrounds MgO particles where CO2 reacts to form CO3 2− ions and H+ ions [34–36]. Free Mg+2 ions could further react with the CO3 2− ions to form MgCO3. However, MgCO3 forms an impervious layer around unreacted MgO particles and hinders the further diffusion of CO2 molecules.

Fig. 2
figure 2

The adsorption/absorption model for MgO–CO2–H2O reaction

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Here, we used mortar and pestle to grind the product. As mentioned earlier, we conducted each experiment for 30 min and ground the sample after that. Grinding helps in scrubbing off the outer nonporous layer of MgCO3. And we performed grinding until we noticed no change in product weight after subsequent experiments. In general, after 3–4 cycles, we observed no change in the weight of product. It is certain that such intermittent grinding step is limited to the laboratory and cannot be seen as an industrial operation. Therefore, it is recommended to have an aggregative fluidization regime for a fluidized bed reactor while scaling up MgO–CO2 reaction. Also, increasing the amount of water vapor cannot lead to the complete carbonate conversion of MgO. Thus, in addition to the amount of steam, surface properties of MgO (such as surface area, particle size, porosity) are also very crucial parameters for the carbonation process.

The thermal analysis curve does not show any significant differences in amounts of carbonate in the high-pressure (50 bars) experiments. We obtained almost similar TGA plots and XRD patterns for different temperatures 300–375 °C and 50 bars CO2 pressure. However, the experiment performed at 300 °C and 10 bars did indicate that the product was not simply MgCO3. The X-ray diffraction pattern in Fig. 1a confirms the presence of Mg (OH) 2 and relatively high amount of MgO·2MgCO3 at 300 °C and 10 bars. DTG curve (Fig. 3a) also evidences the similar situation. The two peaks around 350 and 475 °C were attributed to the losses of water of crystallization and hydroxyl water, respectively. However, another peak at 550 °C corresponds to the complete decomposition of MgCO3. But at a temperature and pressure of 350 °C and 10 bars, respectively, only MgCO3 was formed. This is evidenced by both DTG (Fig. 3b) and X-ray diffraction analysis (Fig. 1a).

Fig. 3
figure 3

DTG plots for product at a 300 °C and 10 bars and b 350 °C and 10 bars

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The steps to calculate % conversion of MgO to MgCO3 is explained in supplementary section [S1]. A conversion of 30.54 and 96.96 % for MgO to MgCO3 was observed at 300 and 350 °C, respectively. Here, we observed the formation of relatively high amount of oxymagnesite (MgO·2MgCO3) at 300 °C and 10 bars. It can be easily calculated that the CO2 capture capacity of oxymagnesite is about two-third that of MgCO3. Hence, formation of oxymagnesite can markedly reduce the overall uptake of CO2 by MgO particles. However, it can be observed from X-ray patterns (Fig. 1) that at higher temperature or pressure, the formation of oxymagnesite is suppressed.

In the same line, recent study reveals that oxymagnesite forms as an intermediate during the thermal decomposition of hydrated magnesium carbonate [37]. Moreover, oxymagnesite can also be formed as a product of reaction between Mg (OH) 2 and CO2 in anhydrous synthesis [29] and/or solid-state reaction of MgO and CO2 using steam [38]. In this regard, Duan et al. [13] calculated the phase diagram of MgO–Mg(OH) 2–MgCO3, which suggests that the transition temperature for direct conversion of MgCO3 to Mg(OH)2 increases with increase in PH2O. Here, when temperature increases from 300 to 350 °C, partial pressure of water also increases and consequently transition temperature also increased. At 300 °C, a relatively less amount of MgO transforms to its carbonates. This can be attributed to a possible high conversion of MgCO3 to Mg(OH) 2, which is in agreement with the phase diagram. Moreover, a significant formation of oxymagnesite at 300 °C and 10 bars also leads to a less direct conversion of MgO to MgCO3. Thus, it can be deduced that a high amount of oxymagnesite forms at a low partial pressure of water. Therefore, the reaction mechanism is changed to:

$${\text{MgO }} + {\text{ CO}}_{ 2} \leftrightarrow {\text{ MgCO}}_{ 3}$$
$${\text{MgO }} + {\text{ 2MgCO}}_{ 3} \leftrightarrow {\text{ MgO}} \cdot 2 {\text{MgCO}}_{ 3}$$
$${\text{MgO}} \cdot 2 {\text{MgCO}}_{ 3} + {\text{ CO}}_{ 2} \leftrightarrow {\text{ 3MgCO}}_{ 3}$$

Moreover, at 50 bars of CO2 pressure and a temperature of 300 °C, MgO has the highest yield of 98.54 %. At high temperatures (300–375 °C) but constant pressure (50 bars), we observe a slight but continuous decrease in the absorption of CO2, which is in congruence with previously reported results [30]. It is well known that at low temperature, the physisorption process dominates but at elevated temperature CO2 chemisorbs on MgO and thus CO2 uptake capacity gradually starts decreasing.

CO2-uptake capacity of a metal oxide is primarily dominated by the factors such as surface area, pore volume, pore functionality and pore size [28]. BET surface area measurement, pore volume and average pore sizes presented in Table 2 indicates the significant role of CO2 to influence the particle structure. As received, 325 mesh size mesoporous MgO has a pore size of 9.09 nm and a high surface area (95.08 ± 1.5 m2/g). Figure 4 illustrates that the N2 adsorption–desorption isotherm curves for as-received MgO follows Type IV isotherms (as per IUPAC classification); typical for mesoporous substances [39]. Also, the hysteresis pattern is H3 (following IUPAC classification) indicating the presence of slit-like pores. After capture of CO2 at 350 °C and 10 bars, surface area dramatically decreased to (4.15 ± 0.05) m2/g. The ratio of surface area to pore volume is also reduced by almost half and results in high diffusion paths. The significant decrease in surface area attributes to the basicity of mesoporous MgO. The mesoporous MgO is highly basic with well-ordered pores to hold high CO2 at both lower and higher temperatures. It is well known that porous materials allow molecules to pass through their pore aperture for storage, separation or conversion [36]. MgO with a fine particle size (<44 µm) has a high content of mesopores, which leads to good mass transfer properties during the absorption process. The mechanism for metal oxide reaction with CO2 gas has been discussed a lot [40–43]. CO2 molecules diffuse through the pores of mesoporous MgO and the present large active sites hold these CO2 molecules [19]. The trapped CO2 molecules further react to form MgCO3.

Table 2 Surface properties of MgO particles before and after carbonation reaction 350 °C, 10 bars
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Fig. 4
figure 4

N2 adsorption–desorption isotherm of as-received MgO

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It can be observed from Table 2 that at 350 °C and 10 bars, 96.96 % of MgO was converted to MgCO3 and almost all the pores have been utilized after reaction with CO2 molecules. Henceforth, almost no CO2 molecules could have diffused further in the pores. Noticeably, pore size increased to 16.25 nm. Thus, it is apparent that the porosity of MgO particles plays a very vital role for CO2 uptake.

Previously, Beruto et al. has reported a very low CO2 absorption capacity of MgO in absence of water vapor [27]. An uptake capacity of 0.089 and 0.091 mmol CO2/g MgO was observed at 350 °C in 1.33 and 3.33 bars of CO2, respectively. Thus, it was concluded that there is a high activation energy barrier to MgO recarbonation. On the contrary, Feng et al. [14] heated the sorbent (MgO) to 1000 °C in pure CO2 and noted a low but almost constant (for 8 cycles) absorption capacity of 0.99 mmol CO2/g MgO. Thus, unlike CaO-based sorbents, MgO does not show a fast decline in their CO2 capture capacity over a large number of carbonation–calcination cycles. Bhagiyalakshmi et al. [20] synthesized basic mesoporous MgO (surface area of 250 m2/g) using mesoporous carbon obtained from SBA-15 and obtained a maximum CO2 adsorption of 2.27 mmol CO2/g at 100 °C and nearly 1.81 mmol CO2/g at 25 °C for a feed flow rate of 30 ml/min CO2 gas (99.9 % purity). Therefore, it is evident that even for a relatively lower surface area (95 m2/g compared to 250 m2/g), MgO can have a better CO2 uptake at a higher temperature (here, 300–375 °C) and higher CO2 pressure (10–50 bars). Recently, Fagerlund et al. [28] reported a very high % conversion of MgO to MgCO3 under high temperature and pressure conditions. However, MgO showed a slow conversion rate (50 % for ~7 h) for the condition [PCO2 (~18–19 bar), PH2O (~1–2 bar) and 300–350 °C] which can be suitable for pre-combustion capture process. The reason could be the deposition of impervious carbonate layer on the surface of MgO which does not allow CO2 molecules to further diffuse in the pores. Therefore, continuous removal of these nonporous layers of carbonates is needed. As mentioned earlier, we recommend to have an aggregative fluidization regime for a fluidized bed reactor while scaling up MgO–CO2 reaction.

In summary, we demonstrate the increase in uptake of CO2 by mesoporous MgO at high temperature and pressure in the humid environment. We observed the significant role of temperature, pressure, water vapor and porosity of the sorbent on the improvement of the CO2 absorption capacity of mesoporous MgO.

Conclusion

MgO is a promising candidate for pre-combustion CO2 capture due to their stable absorption capacity of CO2, adequate mechanical strength, high attrition resistance and good regenerability. We observed a 96.96 % of MgO conversion to MgCO3 at 350 °C and 10 bars CO2 pressure. Moreover, CO2 uptake capacity of MgO can be increased by performing the reaction in the presence of slight humid environment and scrubbing off the outer non-porous layer of MgCO3. We also demonstrate that MgO has a low regeneration temperature (550 °C) and stable CO2 capture capacity over wide range of temperature and pressure. Interestingly, the present study also illustrate that a high amount of oxymagnesite forms at a low partial pressure of water.