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Article

Synthesizing and Characterizing a Mesoporous Silica Adsorbent for Post-Combustion CO2 Capture in a Fixed-Bed System

by
Hind F. Hasan
1,
Farah T. Al-Sudani
1,
Talib M. Albayati
1,*,
Issam K. Salih
2,
Hamed N. Harharah
3,
Hasan Sh. Majdi
2,
Noori M. Cata Saady
4,
Sohrab Zendehboudi
5 and
Abdelfattah Amari
3
1
Department of Chemical Engineering, University of Technology-Iraq, 52 Alsinaa St., Baghdad 35010, Iraq
2
Department of Chemical Engineering and Petroleum Industries, Al-Mustaqbal University, Babylon 51001, Iraq
3
Department of Chemical Engineering, College of Engineering, King Khalid University, Abha 61411, Saudi Arabia
4
Department of Civil Engineering, Memorial University, St. John’s, NL A1B 3X5, Canada
5
Department of Process Engineering, Memorial University, St. John’s, NL A1B 3X5, Canada
*
Author to whom correspondence should be addressed.
Catalysts 2023, 13(9), 1267; https://doi.org/10.3390/catal13091267
Submission received: 25 July 2023 / Revised: 29 August 2023 / Accepted: 30 August 2023 / Published: 2 September 2023 / Corrected: 17 January 2024
(This article belongs to the Topic Materials and Catalysts for Pollutants and CO2 Capture and Transformation)
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Abstract

:
MCM-41, a mesoporous silica with a high surface area and hexagonal structure, was synthesized, and commercial nano-silicon dioxide (SiO2) was used as a solid adsorbed in post-combustion CO2 capture. The CO2 adsorption experiments were conducted in a fixed-bed adsorption system using 5–15 vol.% CO2/N2 at a flow rate of 100 mL/min at varying temperatures (20–80 °C) and atmospheric pressure. Analyses (X-ray diffraction, nitrogen adsorption-desorption isotherms, Fourier-transform infrared spectroscopy, and transmission electron microscopy (TEM)) revealed that the synthesized MCM-41 has mesoporous characteristics: a high surface area and large pore volumes. The CO2 adsorption capacity of MCM-41 and commercial nano-SiO2 increased considerably with increasing CO2 concentration and temperature, peaking at 60 °C. Below 60 °C, dynamics rather than thermodynamics governed the adsorption. Increasing the temperature from 60 to 80 °C decreased the adsorption capacity, and the reaction became thermodynamically dominant. Additionally, compared with commercial nano-SiO2, the MCM-41 sorbent demonstrated superior regenerability and thermal stability.

1. Introduction

Carbon dioxide (CO2) emissions into the atmosphere contribute to global warming and climate change; thus, they pose a significant negative impact on humanity. The effective approach of capturing and storing carbon is essential to limiting global warming to 1.5 °C above preindustrial levels [1,2,3]. Approximately 76% of the total emissions from all carbon sources, weighted by global warming potential (GWP), arise from the combustion of fossil fuels because they are still the principal sources of energy for global consumption [4]. Post-combustion CO2 capture is the technology with the greatest promise for eliminating CO2 because it has the most potential to be adapted to current power-generation units [5]. Following the combustion of fossil fuels (petroleum, natural gas, coal, and diesel) or other carbon-based substance sources (biomass such as wood), CO2 gas is drawn out of the flue gases in the process. Post-combustion CO2 capture can be accomplished using various techniques, such as chemical absorption [6,7], cryogenic separation [8,9], membrane-based separation [10,11], and adsorption separation procedures [1,12]. One of these techniques, amine-based CO2 capture, has been extensively utilized in commercial-scale CO2 capture plants [13,14] because of its advantages. The amine-based CO2 capture process has an affordable operation cost, high efficiency, and can handle a more substantial emission stream. Furthermore, amine-based CO2 adsorbent is highly reactive to CO2, absorbs a small amount of hydrocarbons, and can be regenerated [15,16]. However, this method has several drawbacks. They include high solvent makeup [16,17]; significant energy consumption for absorbent regeneration, accounting for 70% of the operational costs [13,18]; a high level of equipment corrosion; sizable absorber volumes [17]; and degradation of the amine absorbent by SO2, NO2, O2, and metal ions in the flue gases that necessitate a high rate of absorbance [16,17]. Researchers [5,19,20] have been searching for potential alternatives to existing strategies (e.g., absorption, adsorption, and membranes). CO2 adsorption by solid adsorbents is one of the most economical techniques because it avoids the need for expensive sorbent renewal, solvent evaporation, and solvent column corrosion. In this approach, the active substances on the absorbent interact with the CO2 in the exhaust gas to adsorb it physically or chemically. Solid adsorbents with a large surface area, low heat capacity, and considerable stability are often needed to provide a high contact area with CO2, allow effective regeneration, and enable operation at high temperatures [1,5,19]. There are many types of mesoporous silicas, such as silica gel, activated carbon, multi-walled carbon nanotubes, zeolites, and metal-organic frameworks [21,22]. The mesoporous material MCM-41 (Mobil Composition of Matter) has a well-defined and organized structure composed of uniform channels and pore diameters. This structure is attributed to its extremely large surface area and pore volume; thus, MCM-41 is an outstanding option for CO2 adsorption [23,24].
In this study, MCM-41 was synthesized to evaluate its physicochemical properties and ability to adsorb and desorb CO2 from a gas stream (CO2/N2). This study investigated the effects of the bed temperature (20–80 °C) and CO2 concentration (5–15 vol.%) in a CO2/N2 gas mixture on CO2 adsorption. Adsorbent types were also investigated, along with the regeneration of two adsorbents to evaluate the reversibility of CO2 adsorption on the prepared MCM-41 and commercial nanoparticle nano-SiO2 to reduce their replacement costs.

2. Results and Discussion

2.1. Adsorbents Characterization

The XRD pattern for synthesized MCM-41 before adsorption is depicted in Figure 1a, showing a prominent peak at 2θ of 2.0–2.2°, which implies microporosity and the presence of a periodic well-ordered, two-dimensional, hexagonal, long-range order of the channel [1,25,26]. Additionally, it displays two small peaks at 2θ of 3.8–3.6° and 5.8–5.2° distributed to the (d100, d110, and d200) faces of the p6mm structure of the mesoporous MCM-41 [1,26]. The establishment of a hexagonal structure is shown by three peaks (i.e., 100, 110, and 200). Comparing the XRD patterns in Figure 1 (curves a and b, before and after CO2 adsorption, respectively) confirms that the CO2 adsorption did not alter the MCM-41 structure; no change in the XRD pattern or d-spacing values could be observed [1]. The XRD patterns of silica nanoparticles before and after CO2 capture (Figure 2a,b) reveal that the nanoparticle SiO2 was amorphous; there was only a maximum broad scattering centered at 19.8–20.4° that confirmed its amorphous nature [27].
The morphology of the mesoporous materials was analyzed by EDX-SEM. The compositions of MCM-41 and nano-SiO2 were determined by EDX fundamental analysis (Figure 3). The atomic ratio of Si and O is about 1:2, which means that the MCM-41 and nano-SiO2 with high purity revealed a strong affinity for Si and O, confirming silica as the predominant element in the sample.
The SEM images of MCM-41 before and after CO2 adsorption (Figure 3a,b) clearly show the fabrication of the well-ordered, hexagonal structure [1]. A more detailed look at the MCM-41′s surface confirms the presence of mesoporous channels of uniform size and shape with various configurations, including spherical channels, puffy or inflated structures, and smooth surfaces [26,28]. For optimal adsorption efficiency, these swelling shapes are preferred [26]. Panels c and d of Figure 3 depict the SEM images before and after CO2 adsorption via nano-silica, showing amorphous silicate materials with irregular sizes and structures as well as spherical nano-silica agglomerates [1,2,29].
The surface functional groups of the prepared sorbent MCM-41 and commercial nano-silica were evaluated by a Fourier-transform infrared (FTIR) spectrometer. Figure 4 shows the infrared spectra before and after the CO2 adsorption of MCM-41, with bands around 424.34, 447.49, 443.63, 806.25, 869.41, 968.27, 1072.42, 1087.85, 1238, and 1245 cm−1 attributed to the links of the Si–O–Si symmetric and asymmetric stretching vibration, which corresponds to the formation of a condensed mesoporous silica network [30,31,32]. Physically adsorbed H2O caused the 1647 cm−1 band due to its bending vibration. High absorption was observed at 3614.60 cm−1 assigned to free Si–OH groups, whereas a broad absorption band centered at 3595.31 cm−1 was identified to belong to hydrogen-bonded Si–OH groups by physically adsorbed H2O [30,32]. The absence of bands between 2850 and 3000 cm−1, corresponding to the stretching of the –C–H group, indicates that the synthesized material’s pores were successfully purified of the surfactant molecules [33].
The FTIR spectra of nano-SiO2 before and after CO2 adsorption are illustrated in Figure 5. Water molecules in the amorphous SiO2 structure (Si–OH stretching vibration and hydrogen bond) caused the IR band at 3483.44 and 3610.74 cm−1. The bending vibration of H2O molecules (the O–H bending vibration of adsorbed molecular water) accounts for the IR band between 1630 and 1634.35 cm−1 [34,35,36]. A very solid and deep IR band was observed at 1099.43 cm−1, typically associated with TO and LO modes of Si–O–Si asymmetric stretching vibrations. The IR band at 964.41 cm−1 was due to silanol groups; the IR band at 802.39 cm−1 was attributed to Si–O–Si symmetric stretching vibrations; and the IR band at 466.7 cm−1 was caused by O–Si–O bending vibrations [35,37].
Table 1 summarizes the results of measuring the Brunauer–Emmett–Teller (BET) surface area and pore volume of the MCM-41 and nano SiO2. For the MCM-41, a higher surface area was generated (966 m2/g) compared with nano-SiO2 (136.36 m2/g). This is because the hexadecyl trimethyl ammonium bromide (CTAB) cationic surfactant helped to generate more new pores and expand the existing pores as the activating agent. The MCM-41 showed a maximum pore volume of 0.91 cm3/g, followed by nano-SiO2 (0.59 cm3/g). These characteristics allowed for greater adsorbent-adsorbate molecule interaction, increasing CO2 adsorption [38].
Figure 6a shows the thermogravimetric analysis (TGA) profiles of the siliceous MCM-41 under a nitrogen environment, illustrating three separate stages of weight reduction. The initial weight reduction was recorded in the temperature range of 25–170 °C, resulting in a decrease in weight of about 2.7%. This phenomenon can be attributed to the release of physically adsorbed water molecules present on the outer surface of the crystallites, as well as water trapped within the interconnected network of larger and smaller pores found among the aggregated crystallites [39]. Between 170 and 320 °C, it was observed that around 35.05% of the weight losses can be attributed to the decomposition or removal of organic substances [40]. The observed phenomenon occurring between 320 and 550 °C can be attributed to the degradation of residual surfactants and the loss of water resulting from the condensation of neighboring silanol groups, leading to the formation of siloxane bonds. This process accounts for approximately 8.5% of the overall transformation. The sample lost 46.26% of its weight when subjected to temperatures up to 1000 °C. The absence of significant weight loss at 550 °C suggests that the surfactant was effectively eliminated.
Figure 6b represents the weight loss profile with temperature for commercial nano-SiO2. The TGA curve showed a small weight loss at a temperature below 100 °C, which indicates the removal of gases and moisture. Above 150 °C, the TGA curve was maintained, which proved the decomposition temperature of SiO2 is more than 1000 °C.
The TEM images for MCM-41 and commercial nano-SiO2 are depicted in Figure 7. The TEM image of MCM-41 (Figure 7a) reveals the presence of long-range alignment of pores, i.e., a hexagonal array of ordered mesopores, regular channels, and parallel, which is also a direct proof for the transformation from a layered material to a hexagonal phase, which was inferred based on X-rays (Figure 1). The TEM image of commercial nano-SiO2 (Figure 7b) indicates that the presented sample possesses an amorphous structure. In addition, TEM image provides clear visual proof of the presence of pores within spherical silica particles, confirming their nano-structural characteristics.

2.2. CO2 Adsorption Behavior

2.2.1. Effect of Adsorbent Type on the CO2 Adsorption

Figure 8 presents the CO2 breakthrough vs. adsorption time for the MCM-41 and nano-SiO2 adsorbents. The effluent’s CO2 concentration, as indicated by C/C0, steadily rose and approached 1.0 as the adsorbent became saturated. As shown in Figure 8, neither the MCM-41 nor the commercial nanoparticle SiO2 experienced any breakthrough time. Furthermore, the results demonstrated that MCM-41 can adsorb CO2 with an adsorption capacity of 0.63 mmol CO2/g, due to the physisorption mechanism.
Table 2 compares the CO2 adsorption capacities of MCM-41 and the commercial nanoparticle SiO2. The MCM-41 clearly exhibited higher CO2 adsorption than the commercial nano-SiO2 due to its large pore volume and high specific area (Table 1), resulting from the physical interaction between the CO2 and MCM-41 via van der Waals forces or electrostatic force [41]. This performance illustrated that a high specific surface area and extremely microporous materials play an essential part in improved CO2 adsorption [42].

2.2.2. Effect of CO2 Composition on Adsorption Performance

Table 2 and Figure 8 show the experimental breakthrough curves and the estimated adsorption capacity at various initial CO2 concentrations for the MCM-41 and commercial nanoparticle SiO2 adsorbents operated at a fixed feed flow rate (100 mL/min) and bed temperature (200 °C). Table 2 shows that as the CO2 concentration in the gas mixture increased from 5 to 15 vol.%, the breakthrough saturation times decreased, and the adsorbents attained saturation more rapidly. The concentration gradient between the bulk phase and the solid phase is affected by the concentration of CO2 in the gas mixture, which influences the mass transfer zone, significantly [12,43]. This caused enhanced CO2 adsorption, rapid saturation at high CO2 concentrations, and significant adsorption capacity. Monazam et al. [44] and Boonchuay et al. [12] found similar trends.

2.2.3. Effect of the Bed Sorption Temperature on the CO2 Adsorption

The CO2 adsorption was significantly influenced by the adsorption temperature. The MCM-41 and commercial nanoparticle SiO2 exhibited temperature-dependent changes in their CO2 adsorption breakthrough saturation times and adsorption capacities (Table 2).
At 60 °C, the optimum temperature for the MCM-41 adsorbent, the breakthrough saturation time, and adsorption capacity attained their maximum values. The CO2 adsorption efficiency noticeably dropped as the temperature increased from 60 to 80 °C. The general pattern agrees with previously reported results [4,45]. Below 60 °C, the CO2 adsorption process was predominantly driven by dynamics rather than thermodynamics. However, as the temperature was increased, thermodynamics began to predominate. Adsorption is an exothermic process; therefore, it is more efficient at low temperatures [4].
For the reason that physical adsorption is an exothermic reaction caused by the adsorbate bonding to the adsorbent surface walls via van der Waals forces, it does not include any reaction between the adsorbate and adsorbent [46]. The molecular motion kinetics of recently introduced gas adsorbates have improved to be more dynamic, leading to greater diffusion rates over the adsorbent surface in the presence of heat [12,47]. Due to the increased rate of adsorbate diffusion with increasing temperature, the contact between the adsorbate and adsorbent was weakened, leading to a diminished adsorption capacity [12]. At elevated temperatures, commercial nanoparticle SiO2 adsorbents become thermodynamically unfavorable [4] due to a decrease in the breakthrough saturation time and adsorption capacity [48]. The exothermic nature of the physical adsorption process decreases the CO2 adsorption capability at high adsorption temperatures [12,49].

2.2.4. Cyclic CO2 Adsorption

It is necessary to evaluate the reversibility of CO2 adsorption on the MCM-41 and commercial nanoparticle SiO2 to reduce their replacement costs in real applications. Ten cycles were conducted to test the recoverability of both sorbents. The operating conditions were as follows: adsorption temperature (60 °C), simulated flue gas component (10/90 vol.% CO2/N2) with total feed flow rate (100 mL/min), while the desorption condition occurred at 100 °C and a feed flow rate of 40 mL/min N2. Figure 9 illustrates the saturation adsorption capacity following each cycle. The adsorption capacity dropped by 3.31% (from 0.68 to 0.657 mmol-CO2/g-sorbent) after 10 cycles when using the MCM-41. In comparison, the adsorption capacity of the commercial nanoparticle SiO2 declined by 8.42% (from 0.57 to 0.522 mmol-CO2/g-sorbent) after six cycles, demonstrating that the MCM-41 sorbent had superior renewability and proved thermally stable in the cyclic adsorption operation. These advantages suggest that the MCM-41 sorbent is promising for further modifications for practical applications of CO2 capture in the field.

3. Experimental Work

3.1. Materials and Apparatus

A typical MCM-41 synthesis involves four reagents: a solvent (water); a silica precursor (tetraethyl orthosilicate [TEOS]); a cationic surfactant as a template agent (hexadecyl trimethyl ammonium bromide or CTAB); and a catalyst (NaOH). Table 3 provides the properties of the chemicals utilized in this work.

3.2. Preparation of MCM-41

The sol-gel method was employed to synthesize MCM-41, in which TEOS acted as the silica source and CTAB controlled the structure’s development. After dissolving 1 g of CTAB in a mixture of 0.34 g NaOH and 30 mL deionized water, the remaining 5.78 g of TEOS was then added drop by drop over roughly 1 h at ambient room temperature while stirring the mixture. In an autoclave at constant hydrothermal condition (110 °C), the resulting homogenous mixture was crystallized for 96 h. A solid product was filtered and then washed with distilled water to eliminate the remaining surfactants. After that, the solid produced was oven-dried at 40 °C overnight. To generate the white powder MCM-41 (Figure 10), the initial version of MCM-41 was calcined at 550 °C for 6 h, at which time all the surfactant was removed.

3.3. Adsorbent Morphology and Surface Properties

The crystalline structures of the prepared MCM-41 nanocomposite and commercial nano-SiO2 were determined using an X-ray diffraction (XRD) instrument (XRD-6000, Shimadzu Scientific Instruments (SSI), Shimadzu, Kyoto, Japan) with CuKα radiation (λ = 1.5418 A) and a scanning range 2θ of 1–80°. The structure and surface morphology were scanned by an electron microscope with high-resolution scanning electron microscopy (SEM) (Axia™ ChemiSEM™, Thermo Fisher Scientific, Waltham, MA, USA). An energy-dispersive spectrometer (EDX) linked to a scanning electron microscope was used to identify the components of the adsorbents. Transmission electron microscopy (TEM Carl Zeiss-EM10C-100Kv, Oberkochen, Germany) was employed to obtain an image of the interior structure. Fourier-transform infrared spectroscopy (FTIR, Spectrum Two™, PerkinElmer, Waltham, MA, USA) was used to identify the functional group characteristics of the adsorbents. The spectrum was collected from 4000–400 cm−1 after the sample was pulverized and diluted with KBr. The surface properties of the adsorbents were determined via N2 adsorption and desorption on a surface area and pore size analyzer (Q-surf 9600, USA). Before conducting adsorption measurements, the sample was exposed to a 5 h vacuum pretreatment (i.e., 0.5 mmHg) at 250 °C. The specific surface area was determined using the Brunauer–Emmett–Teller (BET) technique. The adsorbed quantity was used to determine the total pore volume after capillary condensation was performed at a relative pressure of P/P0 = 0.994. The Barrett–Joyner–Halenda (BJH) model’s desorption branch was used to estimate the pore size [1,4]. Thermal gravimetric analysis (TGA) (STA PT1000, Linseis, Robbinsville, NJ, USA) was performed for mesoporous silica and commercial nono-SiO2 to evaluate thermal stability (weight loss). The samples (10 mg) were heated from room temperature to 1000 °C at a heating rate of 0 °C/min in an N2 atmosphere with a flow rate of 50 mL/min.

3.4. CO2 Adsorption and Desorption Performance Measurement

3.4.1. CO2 Adsorption Experiments

The CO2 adsorption and desorption experiments were conducted in a fixed-bed quartz reactor. Two types of adsorbents (i.e., MCM-41 and nano-SiO2) were used in the CO2 adsorption at 20, 40, 60, and 80 °C using 5/95, 10/90, and 15/85 vol.% CO2/N2 as a feed gas, indicating the concentration of CO2 in the flue gas from a power plant burning coal. At ambient pressure, 0.5 g of the adsorbent was placed into a quartz tube reactor (0.8 mm ID) in a conventional cycle. To prevent the product from flowing, quartz wool was applied on the side of the gas outlet. Before the CO2 adsorption evaluation, the adsorbent was heated in a tube furnace (BTF-1200-MS, Anhui Bei Keke Equipment Technology Co., Ltd., Anhui, China) at 100 °C with an N2 flow rate of 40 mL/min for 2 h to lower the physically adsorbed moisture and any residual CO2 in the adsorbent. Furthermore, it was cooled to the required experimental temperature. Adsorption of CO2 started at a rate of 100 mL/min after the required percentage of CO2 and the proper temperature were established. The input gas flow rate was regulated using CO2 and N2 flow meters with a range of 0–100 mL/min. The volumetric (bubble) flow meter (Agilent Optiflow Digital Flowmeter, Santa Clara, CA, USA) was utilized to calibrate the flow meters with an accuracy of 3%. A gas analyzer was used to monitor the CO2 concentration at the reactor’s outlet every 20 s until saturation (when the concentration of CO2 in the released gas approximated that of the input gas). Figure 11 displays the experimental setup for the CO2 adsorption-desorption process.
The saturated adsorption capacity of CO2 was displayed as a plot of C/C0 vs. time to establish a breakthrough profile. The integrated equation of the breakthrough curves (Equation (1)) was used to determine the saturation level of the CO2 adsorption capacity for both of the sorbents [4].
Q s = F × ∫ 0 t ( C 0 − C ) d t m
where Qs represents the saturated CO2 adsorption capacity across the MCM-41 and nano-SiO2 adsorbents (mmol/g); F depicts the input CO2 volumetric velocity (mL/min); and m is the adsorbent’s quantity (g). C0 and C stand for the CO2 concentrations in the input and the discharge (mmol/L), respectively; and t is the adsorption period (min). The saturation adsorption capacity (Qs) is defined as the CO2 adsorption capacity of C/C0 = 1.0, while the breakthrough adsorption capacity (Qb) is defined as the CO2 adsorption capacity of C/C0 = 0.05 [4], and the time corresponding to Qb is referred to as the breakthrough time.
The saturated adsorption capacity for CO2 was displayed as a plot of C/C0 vs. time to establish a breakthrough profile. The capacity for saturating sorption of the CO2 over the sorbents (mmol/g) was calculated from the integral of the breakthrough curve using Equation (1) [4].

3.4.2. CO2 Desorption Experiments

The significant adsorption capacity and superior renewability of the absorbent in the cyclic adsorption-desorption process are crucial in practical commercial applications. In this work, 10 cycles were performed to assess the MCM-41 and nano-SiO2 adsorbent’s capacity for regeneration. The experiments related to desorption were continued after adsorption by altering the CO2/N2 feed gas simulation to N2 gas at a flow rate of 40 mL/min and adding heat at 100 °C after the adsorption experiments were completed. This was accomplished by continuously measuring the concentration of CO2 in the discharge flow until the gas analyzer exhibited no more CO2 peaks.
There are numerous research studies with focus on important aspects (e.g., design, performance prediction, thermodynamic behaviors, and economic assessment) of carbon management in the literature [50,51]; however, further research and engineering activities on new strategies such as use of metal organic frameworks (MOFs) as well as efficient and inexpensive chemicals for CO2 separation and capture are still required for optimization and scale-up purposes.

4. Conclusions

This study synthesized and characterized MCM-41 and compared its performance to that of commercial nano-SiO2 for CO2 capture. The synthesized MCM-41 has a well-organized arrangement of mesopores, namely a hexagonal array, characterized by parallel and regular channels. This observation serves as a direct evidence supporting the transition from a layered material to a hexagonal phase, a conclusion previously derived from X-ray analysis. Additionally, commercially, nano-SiO2 exhibits an amorphous structure. The key findings are that the synthesized mesoporous silica MCM-41, which has a large surface area and pore structure, showed a high CO2 adsorption capacity, indicating its high potential for capturing CO2 emissions compared with the nano-SiO2.
Since the saturation adsorption capacity of the MCM-41 and nano-SiO2 declined as the temperature increased, the CO2 adsorption onto these materials was an exothermic reaction according to the physical adsorption affected by van der Waals forces. Furthermore, the results showed that the CO2 adsorption via MCM-41 was dynamic and thermodynamically favorable, while the nano-SiO2 was thermodynamically unfavorable. In addition, after 10 cycles of adsorption-desorption at temperatures ranging from 60 °C (adsorption) to 100 °C (desorption), the MCM-41 showed good renewability, while the nano-SiO2 showed renewability only for six cycles. This makes MCM-41 highly suitable for CO2 capture in the petroleum refinery industry.

Author Contributions

Conceptualization, H.F.H.; methodology, I.K.S.; software, H.N.H. and N.M.C.S.; validation, H.S.M.; formal analysis, H.F.H. and S.Z.; investigation, S.Z.; resources, F.T.A.-S.; data curation, A.A.; writing original draft preparation, T.M.A.; writing and editing, N.M.C.S.; visualization, T.M.A.; supervision, F.T.A.-S.; project administration, H.F.H.; funding acquisition, H.N.H. All authors have read and agreed to the published version of the manuscript.

Funding

This research was funded by the Dean of Scientific Research at King Khalid University under the grant number RGP2/133/44.

Data Availability Statement

All relevant data and material are presented in the main paper.

Acknowledgments

The authors are grateful to the Department of Chemical Engineering at the University of Technology-Iraq, the Department of Chemical and Petroleum Industries Engineering, Al-Mustaqbal University, Babylon, Iraq, and the Department of Civil Engineering, and Department of Process Engineering, Memorial University, St. John’s, NL A1B 3X5, Canada. The authors extend their appreciation to the Deanship of Scientific Research at King Khalid University of funding this work through large group Research Project under grant number RGP2/133/44.

Conflicts of Interest

The authors declare no conflict of interest.

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Catalysts 13 01267 g001
Figure 1. X-ray diffraction patterns of MCM-41: (a)—before and (b)—after CO2 adsorption.
Figure 1. X-ray diffraction patterns of MCM-41: (a)—before and (b)—after CO2 adsorption.
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Figure 2. X-ray diffraction patterns of SiO2: (a)—before and (b)—after CO2 adsorption.
Figure 2. X-ray diffraction patterns of SiO2: (a)—before and (b)—after CO2 adsorption.
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Catalysts 13 01267 g003aCatalysts 13 01267 g003b
Figure 3. EDX—SEM of (a)—Fresh MCM-41, (b)—MCM-41 after adsorption, (c)—Fresh SiO2, and (d)—SiO2 after adsorption.
Figure 3. EDX—SEM of (a)—Fresh MCM-41, (b)—MCM-41 after adsorption, (c)—Fresh SiO2, and (d)—SiO2 after adsorption.
Catalysts 13 01267 g003aCatalysts 13 01267 g003b
Catalysts 13 01267 g004
Figure 4. FT—IR spectra of (a)—Fresh MCM-41, and (b)—MCM-41 after adsorption.
Figure 4. FT—IR spectra of (a)—Fresh MCM-41, and (b)—MCM-41 after adsorption.
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Catalysts 13 01267 g005
Figure 5. FT—IR spectra of SiO2 (a)—Fresh SiO2, and (b)—SiO2 after adsorption.
Figure 5. FT—IR spectra of SiO2 (a)—Fresh SiO2, and (b)—SiO2 after adsorption.
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Figure 6. TGA analysis for (a)—MCM-41 and (b)—nano-SiO2.
Figure 6. TGA analysis for (a)—MCM-41 and (b)—nano-SiO2.
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Catalysts 13 01267 g007
Figure 7. TEM images for (a)—MCM-41 and (b)—nano-SiO2.
Figure 7. TEM images for (a)—MCM-41 and (b)—nano-SiO2.
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Figure 8. Breakthrough profile of CO2 adsorption via (a)—MCM-41 and (b)—Nano-SiO2.
Figure 8. Breakthrough profile of CO2 adsorption via (a)—MCM-41 and (b)—Nano-SiO2.
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Catalysts 13 01267 g009
Figure 9. Cyclic CO2 adsorption capacity of the MCM-41 and commercial nanoparticle SiO2.
Figure 9. Cyclic CO2 adsorption capacity of the MCM-41 and commercial nanoparticle SiO2.
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Figure 10. Synthesis steps for MCM-41.
Figure 10. Synthesis steps for MCM-41.
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Figure 11. Experimental set-up for the CO2 adsorption-desorption.
Figure 11. Experimental set-up for the CO2 adsorption-desorption.
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Table 1. The surface properties of pure MCM-41 and nano-SiO2 after adsorption.
Table 1. The surface properties of pure MCM-41 and nano-SiO2 after adsorption.
Adsorbent Specific Surface Area (BET)
(m2/g)		Total Pore Volume
(cm3/g)
MCM-419960.91
Nano-SiO2136.360.59
Table 2. CO2 adsorption capacities and saturation time of the MCM-41 and nano-SiO2.
Table 2. CO2 adsorption capacities and saturation time of the MCM-41 and nano-SiO2.
Bed Temp. (°C)Adsorption Capacity
(mmol-CO2/g-Sorbent)		Saturation Time
(min)
MCM-41
5%10%15%5%10%15%
200.480.560.614.52.51.8
400.550.620.7152.72.3
600.620.680.73632.3
800.50.620.724.82.82.3
Nano-SiO2
5%10%15%5%10%15%
200.60.660.714.52.52.5
400.530.600.6752.52.5
600.450.570.6163.52.5
800.310.430.526.33.82.9
Table 3. The chemical properties utilized in the experiment.
Table 3. The chemical properties utilized in the experiment.
Chemical CompoundChemical
Formula 		Molecular Weight
(g/mole)		OriginPurity
(wt.%)
Tetraethyl orthosilicate (TEOS)Si (OC2H5)4208.33Hubei Bluesky New Material Inc., Xiantao, China98%
CetyltrimethylammoniumBromide (CTAB)C19H42BrN364.45Interchiniques SA, France98%
Sodium hydroxideNaOH39.99Didactic, Barcelona, Spain99%
Nano-Silicon dioxideSiO260.08Hubei Bluesky New MaterialInc., Xiantao, China>98%
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Hasan, H.F.; Al-Sudani, F.T.; Albayati, T.M.; Salih, I.K.; Harharah, H.N.; Majdi, H.S.; Saady, N.M.C.; Zendehboudi, S.; Amari, A. Synthesizing and Characterizing a Mesoporous Silica Adsorbent for Post-Combustion CO2 Capture in a Fixed-Bed System. Catalysts 2023, 13, 1267. https://doi.org/10.3390/catal13091267

AMA Style

Hasan HF, Al-Sudani FT, Albayati TM, Salih IK, Harharah HN, Majdi HS, Saady NMC, Zendehboudi S, Amari A. Synthesizing and Characterizing a Mesoporous Silica Adsorbent for Post-Combustion CO2 Capture in a Fixed-Bed System. Catalysts. 2023; 13(9):1267. https://doi.org/10.3390/catal13091267

Chicago/Turabian Style

Hasan, Hind F., Farah T. Al-Sudani, Talib M. Albayati, Issam K. Salih, Hamed N. Harharah, Hasan Sh. Majdi, Noori M. Cata Saady, Sohrab Zendehboudi, and Abdelfattah Amari. 2023. "Synthesizing and Characterizing a Mesoporous Silica Adsorbent for Post-Combustion CO2 Capture in a Fixed-Bed System" Catalysts 13, no. 9: 1267. https://doi.org/10.3390/catal13091267

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