Cite this article as: | Chenguang Qian, Chunquan Li, Peng Huang, Jialin Liang, Xin Zhang, Jifa Wang, Jianbing Wang, and Zhiming Sun, Research progress of CO2 capture and mineralization based on natural minerals, Int. J. Miner. Metall. Mater., 31(2024), No. 6, pp.1208-1227. https://dx.doi.org/10.1007/s12613-023-2785-4 |
Over the centuries, human actions have resulted in a continual increase in the average atmospheric concentration of CO2, which rose from less than 280 ppm before the Industrial Revolution to 418 ppm in 2022 (Fig. 1(a)) [1–2]. In general, CO2 emissions are caused by the burning of fossil fuels, such as coal, oil, and natural gas (Fig. 1(b)) [2]. CO2 concentration has consistently increased by approximately 2.3 ppm/year in the previous decade. If CO2 emissions remain unregulated, the concentration of CO2 in the atmosphere will exceed 500 ppm by 2050 [3]. Considerable CO2 emissions result in global warming, which in turn triggers a host of issues, including the melting of glaciers, acidification of oceans, rise in sea levels, extinction of various animal and plant species, frequent extreme weather, and rampant disease [4–7]. Global warming has a notable influence on social, economic, and human survival and development [8].
Many countries are committed to the development of hydropower, nuclear power, wind power, hydrogen, and other clean energy. The energy structure dominated by fossil fuels, such as coal, oil, and natural gas, will remain unchanged for a long time in the future, and the avoidance of considerable amounts of CO2 emissions will be difficult [9–10]. In this context, carbon capture utilization and storage (CCUS/CCS) technologies have been proposed and used to mitigate CO2 emissions [11–18]. CCUS/CCS technology involves the separation of CO2 emissions from the industry and related energy industries through carbon capture technology. The captured CO2 is then either stored or utilized to reduce CO2 emission (Fig. 2) [19]. CCUS/CCS technology can achieve widespread low-carbon utilization of fossil energy, and it has attracted great attention from the international community. In addition, CCUS/CCS technology is the most direct and effective measure for the control of CO2 emissions in the world and will be widely used in power generation and other industries in the future [20–22]. By 2050, 7 billion tons of CO2 is expected to be sequestered annually [23].
Carbon capture technology separates CO2 from industrial waste gas to obtain high-purity CO2. Based on the different coupling positions in the combustion process, capture methods are divided into three categories: pre-combustion capture, post-combustion capture, and oxygen-enriched combustion (Fig. 3) [12,23–24]. Pre-combustion and oxyfuel combustion contains high concentrations of CO2, which can be directly used or stored. The flue gas obtained after combustion has the characteristics of low pressure, low CO2 concentration, and high N2 content. Thus, high-concentration CO2 can be obtained by carbon capture technology [25].
Flue gas after combustion can be captured using methods such as absorption, solid adsorption, membrane separation, cryogenic separation, and hypergravity. Absorption occurs as physical and chemical absorption [26–27]. Physical absorption is dependent on water, methanol, propylene carbonate, and other solutions as absorbents. According to Henry’s law, the solubility of CO2 in solution changes with the pressure to absorb or desorb [28]. Chemical absorption refers to the absorption of CO2 through acid–base reaction [14]. Membrane separation separates specific gases from mixed ones by utilizing the pressure difference on both sides of a membrane as the separation power or permeability difference of certain polymer membranes to distinguish gases [29]. Solid adsorption utilizes the van der Waals force between CO2 molecules and adsorbents to adsorb and desorb CO2 under various temperatures and pressure conditions. This method has the advantages of low energy consumption and corrosion, simple process, high stability, and environmental protection [30–31]. Given the adsorption–desorption temperature of solid materials, solid adsorbents can be divided into low, medium, and high-temperature capture materials [2]. The capture materials in the low-temperature region (<473 K) mainly include metal–organic frameworks (MOFs), zeolites, porous carbon materials, mesoporous silica, porous organic polymers, and natural minerals [25,32–35]. The middle-temperature region (473–673 K) mainly contains capture materials, including metal oxides and hydrotalcite-like compounds [36]. The high-temperature region (>673 K) mainly comprises capture materials including lithium zirconate, natural minerals with high calcium and magnesium content, and industrial solid waste (Fig. 4) [2,37–39]. These solid adsorbents can be utilized in CO2 capture at different temperatures. Natural minerals are exceptional solid adsorption matrix materials due to their plentiful reserves, low cost, high specific surface area, rich pore structure, excellent mechanical properties, and chemical stability [40].
Rather than releasing it into the atmosphere, carbon sequestration technology captures, compresses, and transports the CO2 generated by large emission sources to select sites for long-term storage. CO2 storage methods mainly consist of geological storage, marine storage, mineral storage (also known as mineralization), etc. [41]. Geological and marine storage are effective methods for CO2 storage; however, these come with risks of carbon dioxide leakage, groundwater pollution, geological disaster, seawater acidification, ecological imbalance, and other risks in the storage process [42–43]; notably, mineralization overcomes these shortcomings [20]. CO2 mineralization refers to the simulation and acceleration of the weathering process of rocks in nature, with CO2 used to react with natural minerals containing calcium and magnesium and form stable carbonates (CaCO3/MgCO3) (Fig. 5) [42,44]. The formation of stable carbonates through CO2 mineralization prevents the harm caused by CO2 leakage in the later stage. Mineralization shows better stability than other storage methods, and the mineralized products have a certain recycling value [45]. In addition, natural minerals, which have gradually become the focus of current research, play a crucial role in CO2 capture and mineralization [46]. This paper reviews the research on CO2 capture and mineralization by natural minerals. Various methods for enhancing the CO2 capture and mineralization performance of natural minerals are also introduced.
Natural minerals play an important role in CO2 capture, utilization, and mineralization due to their plentiful reserves, low cost, excellent mechanical properties, and chemical stability. CO2 capture frequently involves the use of minerals, such as kaolinite, halloysite, montmorillonite, bentonite, attapulgite, sepiolite, etc. The capture of CO2 by these minerals primarily occurs through physical adsorption; however, their adsorption capacity is relatively limited [47]. Therefore, the CO2 capture capability of natural minerals must be improved via different modification strategies [25].
In general, CO2 adsorption onto the surface and interior of natural minerals occurs via physical adsorption or chemical reactions. The adsorption capacity of CO2 is influenced by crystal structure, specific surface area, pore structure, ion exchange, and other properties of minerals. Table 1 shows the CO2 adsorption capacity of natural minerals.
Support | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||
Temperature / °C | Gas atmosphere | Pressure / MPa | |||
Kaolinite | 0 | CO2 | 3 | 0.29 | [48] |
25 | 3 | 0.14 | [48] | ||
25 | 0.1 | 0.06 | [49] | ||
15 | 1.7 | 0.14 | [50] | ||
Halloysite | 0 | CO2 | 3 | 6.17 | [51] |
75 | 0.1 | 0.08 | [52] | ||
Bentonite | 25 | CO2 | 0.1 | 0.13 | [53] |
30 | 0.14 | [54] | |||
Montmorillonite | 25 | CO2 | 0.1 | 0.16 | [55] |
45 | 0.22 | [56] | |||
Palygorskite | 45 | CO2 | 0.1 | 0.27 | [56] |
25 | 0.40 | [57] | |||
Sepiolite | 45 | CO2 | 0.1 | 0.93 | [56] |
50 | 0.41 | [58] |
The mineral known as kaolinite (Al4[Si4O10](OH)8) is a typical 1:1 octahedral layered clay mineral consisting of a SiO4 tetrahedron layer and an AlO2(OH)4 layer [49]. Strong hydrogen bonds tightly connect the interlayer, resulting in low CO2 adsorption capacity due to the dominant physical adsorption [59]. Meanwhile, halloysite (Al2Si2O5(OH)4·nH2O) is a polytype of kaolinite with a multiwalled nanotubular structure. The outer surface comprises tetrahedral groups containing silicon–oxygen groups (Si–O–Si), and the inner wall consists of octahedral aluminum hydroxyl groups (Al–OH) [60–61]. This mineral has high specific surface area, developed pore structure, and rich active groups, which are conducive to CO2 adsorption. Montmorillonite ((Na,Ca)0.33(Al,Mg)2[Si4O10](OH)2·nH2O) is a layered aluminosilicate with a 2:1 structure, and it consists of two silicon–oxygen tetrahedron layers and one aluminum–oxygen octahedron layer. Compared with that of kaolinite, the interlayer of montmorillonite is more prone to relative sliding. Montmorillonite exhibits a natural nanosheet morphology, high porosity, and good cation exchange, and it has a slightly higher CO2 adsorption capability than kaolinite [62]. Attapulgite (Mg5Si8O20(OH)2(OH2)4·4H2O) or palygorskite recognized through its 2:1 layered chain structure composed of a silicon–oxygen tetrahedral layer and a discontinuously arranged octahedral layer. This natural magnesium–aluminum-rich clay mineral is known for its rich pores, large specific surface area, and high porosity, which can be attributed to its unique fibrous, chain, rod-like microstructure, and nanoproperties. Moreover, attapulgite contains numerous oxygen-containing functional groups, which are evenly distributed on its surface and enhance the interaction between adsorbents and CO2 [63]. Sepiolite (Si12Mg8O30(OH)4(OH2)4·8H2O) is a lightweight hydrated magnesium silicate clay mineral. This mineral possesses a nanofiber structure and a high number of silane alcohol groups, uniform pore size, high porosity, and large specific surface area, which can enhance CO2 adsorption [64].
To enhance the adsorption capability of natural minerals to CO2, scientists treat them by using common treatment methods, including pretreatment (heat, acid, alkali, etc.), organic amine, ionic liquid (IL), pillared or metal-doped modification, etc.
In the production of solid CO2 adsorbents, natural minerals must first be treated with heat, acids, and alkali, to further improve their purity, increase their specific surface area, and enhance their pore structure to increase the number of CO2 adsorption sites and reaction cavities. Table 2 reveals the changes in structural and adsorption properties of materials after the treatment of natural minerals by heat, acids, and alkali.
Supports | Pretreatment | Structure property | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||||
Specific surface area / (m2·g−1) | Pore volume / (cm3·g−1) | Temperature / °C | Gas atmosphere | Pressure / MPa | |||||
Kaolinite | Without pretreatment | 8 | 0.01 | — | — | — | — | [65] | |
6 M HCl, 4 h | 38 | 0.07 | |||||||
5 M NaOH, 4 h | 16 | 0.03 | |||||||
Halloysite | Without pretreatment | 63.4 | 0.24 | — | — | — | — | [66] | |
850°C, 4 h + 6 M HCl, 6 h | 366.4 | 0.55 | |||||||
Bentonite | Without pretreatment | 81 | 0.10 | 30 | CO2 | 0.1 | 0.14 | [54] | |
3 M HCl, 3 h | 227 | 0.31 | 0.34 | ||||||
Montmorillonite | Without pretreatment | 72 | 0.16 | — | — | — | — | [65] | |
6 M HCl, 4 h | 253 | 0.71 | |||||||
5 M NaOH, 4 h | 183 | 0.56 | |||||||
Sepiolite | Without pretreatment | 79 | 0.01 | 50 | CO2 | 0.1 | 0.41 | [58] | |
2 M HCl, 24 h | 182 | 0.21 | 0.65 | ||||||
Without pretreatment | 42.74 | 0.08 | 75 | CO2 | 0.1 | 0.27 | [64] | ||
2 M HCl, 6 h | 168.96 | 0.14 | 0.41 | ||||||
Without pretreatment | 42.74 | — | 70 | CO2 | 0.1 | 0.30 | [67] | ||
4 M HCl, 6 h | 320.31 | 0.50 |
Heat treatment increases a material’s specific surface area and pore structure [68]. Its treatment effect mainly depends on the temperature and adsorbate properties. This process can be used as a pretreatment step in combination with acid or alkali treatment. Niu et al. [66] first calcined halloysite in air at 850°C for 4 h and added the calcined halloysite with 6 mol/L hydrochloric acid; the specific surface area of the treated halloysite increased from 63.4 to 366.4 m2/g, which can provide more organic amine loading sites. Acid treatment (e.g., HCl, H2SO4, and HNO3) can be used to leach metal cations (such as Al3+, Mg2+, Fe3+, Na+, etc.) and impurities in natural minerals, considerably increase the specific surface area and pore volume, and improve structural properties, thereby increasing the adsorption performance of CO2 [47]. Zhu et al. [58] applied hydrochloric acid treatment of sepiolite; as a result, hydrogen ions replaced magnesium ions, and a large amount of Si–OH was produced on the pore wall of the mineral. An increase in the specific surface area was observed from 79 m2/g before acid treatment to 182 m2/g afterward. The adsorption capacity of CO2, mainly physical adsorption, increased from 0.41 to 0.65 mmol/g. Ouyang et al. [67] conducted the pretreatment of sepiolite using 4 mol/L hydrochloric acid; they observed that the specific surface area of the treated sepiolite reached up to 320 m2/g, which was eight times higher than that of pristine sepiolite; moreover, the adsorption of CO2 was enhanced from 0.30 to 0.50 mmol/g at 70°C. Compared with acid treatment, metal cations may be inert in alkali treatment (such as NaOH), which results in the poor modification effect of natural minerals [65]. In addition, appropriate heat treatment temperature and acid/alkali concentration are required in the pretreatment of natural minerals using heat, acid, alkali, and other methods. Excessive acid and alkali concentrations and high heat treatment temperatures may damage the structure of natural minerals, thus affecting CO2 adsorption properties.
Organic amine absorbents enable the chemical interaction between amine groups in molecules and CO2. Therefore, the application of these absorbents in CO2 capture is one of the most effective methods for enhancing CO2 capture in the industry [69–70]. The commonly used organic amine absorbents include amine alcohols (monoethanolamine (MEA), diethanolamine (DEA), and methyl diethanolamine (MDEA), aliphatic amines (ethylenediamine (EDA), tetraethylenepentamine (TEPA), hexylamine (HXA), dodecylamine (DDA), and octadecyl amine (ODA)), amides (pentavinylacetamide (PEHA) and formamides), and other amines (polyethyleneimine (PEI)). However, as a result of equipment corrosion, solvent degradation, and high energy-intensive regeneration during the use of organic amine absorbents, their application in the industry faces enormous challenges [71–73]. Researchers have proposed the inclusion of amine-based absorbents with an affinity for CO2 into the structures of natural minerals through the physical leaching method to increase the mass transfer and rapid adsorption of CO2 [48,64,66,74–75] (Fig. 6); this method improves the CO2 capture capacity of amine-based composite solid adsorbents and reduces energy consumption during the adsorbent regeneration process [76]. The CO2 adsorption capacity of organic amine-functionalized natural minerals can be found in Table 3.
Support | Pretreatment | Amine loading / wt% | Adsorption conditions | Regeneration conditions | Cycle number |
CO2 uptake / (mmol·g−1) | Loss / % |
Ref. | |||||
Temperature / °C |
Atmosphere | Pressure / MPa |
Temperature / °C |
Atmosphere | X | Xresidue | |||||||
Kaolinite | Intercalation | 6.35 (HXA) | 0 | CO2 | 3 | 105 | N2 | — | 0.75 | — | — | [48] | |
12.6 (DDA) | 0.50 | ||||||||||||
41.7 (ODA) | 0.28 | ||||||||||||
Without pretreatment | 20 (MEA) | 25 | CO2 | 0.1 | 100 | N2 | — | 2.94 | — | — | [49] | ||
33.3 (EDA) | 1.35 | ||||||||||||
50 (4MEA–1EDA) | 3.38 | ||||||||||||
Halloysite | Without pretreatment | 44.4 (Formamide) | 0 | CO2 | 3 | 250 | — | — | 6.1 | — | — | [51] | |
25 | 3.4 | ||||||||||||
Without pretreatment | 30 (PEI) | 75 | CO2 | 0.1 | 75 | N2 | 10 | 1.5 | 1.4 | 6.6 | [52] | ||
Hydrochloric acid treatment | 50 (PEI) | 85 | 40vol%N2 + 60vol%CO2 | 0.1 | 110 | N2 | 10 | 2.75 | 2.66 | 3.2 | [66] | ||
Hydrochloric acid treatment | 30 (PEI) | 25 | Dry air | 0.1 | 80 | — | 50 | 1.24 | 1.20 | — | [77] | ||
Palygorskite | Hydration treatment | 37 (PEI) | 45 | CO2 | 0.1 | 110 | Ar | — | 1.52 | — | — | [56] | |
Sepiolite | Hydrochloric acid treatment | 47 (PEHA) | 50 | CO2 | 0.1 | 105 | N2 | 10 | 2.47 | 2.23 | 9.8 | [58] | |
15vol%CO2 + 85vol%N2 (water content 5vol%) | — | 2.94 | — | — | |||||||||
Hydrochloric acid treatment | 50 (PEI) | 75 | 40vol%N2 + 60vol%CO2 | 0.1 | 110 | N2 | 10 | 2.48 | 2.31 | 6.8 | [64] | ||
Hydrochloric acid treatment | 60 (TEPA) | 60 | 1vol%CO2 + 99vol%N2 (water content 1vol%) | 0.1 | 90 | N2 | 10 | 3.8 | 3.5 | 7.9 | [78] | ||
Bentonite | Without pretreatment | 30 (PEI) | 75 | CO2 | 0.1 | 75 | He | 5 | 1.06 | 1.02 | 3.7 | [53] | |
Sulfuric acid treatment | 50 (TEPA) | 75 | CO2 | 0.1 | 100 | N2 | 10 | 2.54 | 2.43 | 4.3 | [79] | ||
15vol%CO2 + 4.5vol%O2 + 80.5vol%N2 (water content 18vol%) | — | 4.31 | — | — | |||||||||
Montmorillonite | Hydration treatment | 37 (PEI) | 45 | CO2 | 0.1 | 110 | Ar | — | 1.03 | — | — | [56] | |
Pillaring and calcination | 60 (TEPA) | 25 | CO2 | 0.1 | — | — | — | 1.64 | — | — | [80] | ||
60 (PEI) | 1.45 | ||||||||||||
Hydrochloric acid treatment | 50 (PEI) | 75 | CO2 | 0.1 | 100 | N2 | 10 | 2.54 | 2.43 | 4.4 | [65] | ||
15vol%CO2 + 4.5vol%O2 + 80.5vol%N2 (water content 3vol%) | — | 3.23 | — | — | |||||||||
Without pretreatment | 20–30 (ODA) | 25 | CO2 | 5 | — | — | — | 7.16 | — | — | [81] | ||
Note: X—CO2 uptake for fresh sample; Xresidue—CO2 uptake for cycle sample. |
For the enhanced adsorption capacity of aminated mineral-based solid adsorbents, the effective adsorption area of minerals and the active sites per unit area must be increased, and appropriate adsorption conditions must be selected [25]. The specific surface area and porosity of solid adsorbents depend on the type and pretreatment method of natural minerals, thus affecting the adsorption area. To a certain extent, the number of active sites of adsorbents depends on the type of organic amine, length of alkyl amine molecular chain, and loading amount. The sufficient contact and reaction between CO2 and aminated solid adsorbents can be guaranteed by appropriate adsorption conditions (such as CO2 flow rate, adsorption temperature, pressure, and water content).
The type of organic amine greatly influences the CO2 adsorption performance. Chen and Lu [49] explored the effect of different amine-modified kaolinite on the adsorption performance of CO2. MEA, EDA, and 4MEA + 1EDA were loaded on kaolinite via immersion method, and adsorption was carried out at the adsorption temperature of 25°C and CO2 flow rate of 30 mL/min. The results show that the CO2 adsorption capacities of MEA, EDA, and 4MEA + 1EDA-modified kaolinite samples reached 2.94, 1.35, and 3.38 mmol/g, respectively. The chain length of alkylamine molecules also plays an important role in CO2 adsorption. A small proportion of amine groups reacts with CO2 in long-chain alkylamine molecules, whereas long-chain alkyl molecules may block the pores of adsorbents, which is not conducive to CO2 diffusion. Liu et al. [48] explored the effect of molecular chain length on the CO2 adsorption performance of intercalated kaolinite. First, kaolinite was intercalated with dimethyl sulfoxide (DMSO) to prepare K/DMSO, which was then dispersed in an HNO3/methanol (MeOH) solution. After heating and washing, K/MeOH was prepared. Finally, HXA, DDA, and ODA were inserted into the modified kaolinite to improve the utilization rate of organic amines, and K/HXA, K/DDA, and K/ODA were prepared as adsorption materials (Fig. 6(a)). The results show the negative correlation of the adsorption capacity of the prepared adsorbent for pure CO2 with the length of the alkylamine molecular chain (K/HXA > K/DDA > K/ODA); this finding indicates that the adsorption of kaolinite alkylamine on CO2 was not simple physical adsorption, and NH2 in the alkylamine molecule played a decisive role in the adsorption process. Nevertheless, excessive amine loading leads to mineral pore blockage and hinders CO2 adsorption. Niu et al. [66] prepared a novel mesoporous silica nanotubes (MSiNTs)/PEI nanocomposite for CO2 capture by impregnating acid-treated halloysite nanotubes (HNTs) with PEI (Fig. 6(c)). A thermogravimetric analyzer was used to analyze the effect of PEI loading on the adsorption capacity of the composite. The results showed that when the adsorption temperature was 50°C, the PEI loading was in the range of 30wt%–50wt%, and the adsorption capacity increased between 1.29–1.83 mmol/g. However, the CO2 adsorption capacity decreased when the PEI loading reached 60wt%. Adsorption temperature and CO2 pressure are also key factors affecting the adsorption capacity of materials. The adsorption temperature is determined based on the state of CO2 to be adsorbed. Given the exothermic nature of the CO2 adsorption process, the amount of CO2 adsorbed at high temperatures is reduced [82]. Ramadass et al. [51] explored the effect of formamide-intercalated halloysite on the adsorption capacity for pure CO2 at 0, 10, and 25°C, and a pressure of 3 MPa. The results indicate that the adsorption capacity of the intercalated halloysite adsorbent was 3.4 mmol/g at 25°C, and it increased to 6.1 mmol/g at 0°C. During CO2 adsorption under normal pressure, an appropriate adsorption temperature benefits the improvement of amine activity and causes the appearance of amine active sites, which can also increase the diffusion of CO2 in the pores, thereby improving the adsorption performance of the material. However, at extremely high adsorption temperatures, thermodynamics becomes dominant, and the equilibrium shifts to the desorption direction of CO2, resulting in a decreased adsorption capacity of CO2. Wang et al. [65] investigated the effect of PEI-impregnated montmorillonite on the adsorption performance of pure CO2 at different adsorption temperatures of 30, 45, 60, 70, 80, and 85°C. The results show that the adsorption capacity of the material to CO2 increased with the increase in adsorption temperature. The sample had a maximum adsorption capacity of 2.54 mmol/g at 75°C, but its adsorption capacity decreased when the adsorption temperature was continuously increased from 75 to 85°C. To simulate the performance of adsorbents in CO2 capture in the actual industry, Zhu et al. [58] designed a fixed-bed reaction system (Fig. 7(a)) and used PEHA-impregnated acid-modified sepiolite as an adsorption material. Different CO2 flow rates affect the CO2 adsorption performance (Fig. 7(b)). An extremely low flow rate is not conducive to CO2 diffusion in adsorbent materials, and an extremely high flow rate leads to transfer resistance between CO2 molecules and active sites, which is not advantageous to the retention of CO2 molecules in adsorbent materials. At a flow rate of 30 mL/min and an adsorption temperature of 50°C, experiments were conducted to determine the effect of water content on the adsorption performance of CO2 for 15vol% CO2 and 85vol% N2 mixed gas. Studies have shown that an appropriate amount of water benefits the adsorption of CO2, but when the water content exceeds 5vol%, the CO2 adsorption reaction is inhibited (Fig. 7(c)). This phenomenon can be attributed to the excessive water molecules covering the pores and active sites of amines, and as a result, CO2 molecules cannot bind to some amine sites, resulting in a decreased adsorption performance of the material.
Irani et al. [78] explored the effect of TEPA-modified sepiolite on CO2 adsorption capacity under water-containing conditions. Studies have shown that the addition of 1vol% water benefits CO2 adsorption, and the reaction mechanism of CO2 under dry and wet conditions varies (Fig. 8). Under wet conditions, two bicarbonates can form between one CO2 molecule and one amino group, and carbamate can form between two amino groups and one CO2. However, only carbamate has been formed under dry CO2 conditions [65,83].
Aminosilane grafting, which is more stable than organic amine leaching, is carried out using a chemical reaction between a silanol group on the mineral surface and an amine–alkoxysilane compound [54,84–85] (Fig. 9)). Stevens et al. [86] prepared diamine-modified montmorillonite through cetyltrimethylammonium bromide (CTAB)-assisted exfoliation and aminopropyltrimethoxysilane (AEAPTS) grafting. The results reveal that the adsorption capacity of CO2 reached 1.8 mmol/g at the CO2 concentration of 15vol% and adsorption temperature of 95°C. However, amino silane grafting exhibited a lower loading amount than organic amine leaching, resulting in a low adsorption capacity. Table 4 presents the CO2 adsorption capacity of amino silane-grafted natural minerals.
Support | Pretreatment | Modifier | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||
Temperature /°C | Atmosphere | Pressure / MPa | |||||
Halloysite | Microwave-assisted acid treatment | APTES | 25 | CO2 | 0.1 | 0.11 | [52] |
Palygorskite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.86 | [56] |
DT | 1.29 | ||||||
Microwave-assisted acid treatment | APTES | 25 | CO2 | 0.1 | 0.75 | [57] | |
Calcination and hydrochloric acid treatment | APTES | 40 | Air | 0.1 | 1.81 | [84] | |
Sepiolite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.99 | [56] |
DT | 1.39 | ||||||
Microwave-assisted acid | APTES | 25 | CO2 | 0.1 | 1.0 | [57] | |
Montmorillonite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.75 | [56] |
DT | 1.13 | ||||||
Inserted layer treatment | AEAPTS | 95 | 15vol%CO2 +85vol%N2 | 0.1 | 1.80 | [86] | |
Bentonite | Hydrochloric acid treatment | APTES | 30 | CO2 | 0.1 | 0.66 | [54] |
Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.98 | [56] | |
DT | 0.73 | ||||||
NaCl and CTAB treatment | APTES | 25 | CO2 | 0.1 | 1.14 | [80] | |
Note: DT represents N1-(3-trimethoxysilylpropyl) diethylenetriamine. |
When amino silane is used to modify different minerals, especially fibrous clay minerals, without pretreatment (such as sepiolite and palygorskite), it poses other modification effects that may not necessarily increase the adsorption capacity for CO2. Such a condition is due to the likelihood of amino silane blocking the nanopore structure within minerals during grafting, which is not conducive to CO2 diffusion [56–57]. Cecilia et al. [57] investigated the effect of APTES-grafted acid-modified sepiolite and palygorskite on CO2 adsorption performance. Studies have shown that after APTES grafting, the pore structure of sepiolite and palygorskite is blocked, which is not beneficial to the physical adsorption of CO2 and thus results in a decreased CO2 adsorption performance.
To further enhance the adsorption capacity of amino silane-grafted minerals and fully take advantage of the natural properties of minerals, some scholars have proposed combining minerals with other porous materials to prepare porous composite materials. Park et al. [85] first prepared HKUST-1 MOF@HNT composites through vacuum loading and solvothermal reaction and HKUST-1@HNT composites via internal corrosion and external surface amination using sulfuric acid and APTES (Fig. 9(c)). Studies have shown that HNT acts as a nanocarrier, the internal HKUST-1 improves gas adsorption capacity, and the grafting of amino silane on the HNT outer surface promotes CO2 adsorption selectivity.
Traditional organic amine leaching solid adsorbents suffer from insufficient stability during cyclic regeneration. On the one hand, they are prone to aging during cycling because of the low boiling point of organic amines. On the other hand, organic amines exhibit a weak binding capability toward the carrier surface, which may cause the detachment of amines during the adsorption process [87]. Although amino silane-grafted solid adsorbents are more stable than organic amine-impregnated solid adsorbents, the adsorption capacity of modified solid adsorbents is limited. Therefore, new solid adsorption materials must be explored.
ILs have attracted extensive attention due to their high capacity, good chemical and thermal stability, easy separation, and recyclability. To improve the utilization efficiency of ILs, reduce costs, and improve the cycle performance of materials, Chen et al. [87] used an ion-exchange method to synthesize an amine polycarboxylate IL [P4442]2[IDA] with multiple adsorption sites for the modification of attapulgite and synthesized a new solid adsorption material for CO2 capture. An adsorption capacity of 1.531 mmol/g was observed at the adsorption temperature of 30°C and [P4442]2[IDA] loading of 30wt%. After 10 cycles, the adsorption capacity of the solid material for CO2 was reduced slightly, which indicates an excellent recycling performance. Hu and Zhang [88] prepared adsorption materials using halloysite as a nanomineral adsorbent matrix and acetate–ethanolamine solution as a modifier. The results show that the adsorption efficiency of halloysite modified by IL can reach 23.33% at room temperature and atmospheric pressure, and such a value is considerably higher than that of CO2 adsorption after calcination.
Pillared interlayer clay (PILC) is a two-dimensional molecular sieve porous material known for its large specific surface area, uniform pore structure, and large interlayer spacing. This material shows promise as a catalyst and adsorption material [89–90]. In addition, as CO2 is an acidic gas, some alkaline metals can be doped into minerals. Alkaline metals contain active sites that can interact with CO2, which enhances the adsorption capacity of pillared clays for CO2 [91–92]. Wang et al. [22] explored the effects of Al2O3, ZrO2, and TiO2–SiO2 pillared montmorillonites on the CO2 adsorption performance. The CO2 adsorption performance of different pillared montmorillonites increased with the increase in pore volume and specific surface area. The TiO2–SiO2 pillared montmorillonite exhibited the largest specific surface area and pore volume and the best CO2 adsorption capacity of 1.18 mmol/g at 273 K and 1 atm. Wu et al. [92] studied the effect of alkali metal (Li, Na, K, and Cs) doping on the CO2 adsorption performance of Al-pillared montmorillonite. In the preparation of alkali metal composite materials, the main method involved the addition of aluminum pillared montmorillonite to metal nitrates and pretreatment under ultrasonic radiation at 60°C. Then, calcination of the dried sample in air was conducted at 400°C for 12 h to obtain the composite material. Studies have shown that aluminum pillaring changes the specific surface area of montmorillonite and enhances its physical adsorption performance. Meanwhile, alkali metal doping mainly increased the alkalinity of samples and improved their chemical adsorption performance. The CO2 adsorption capacities of all samples had the following order 5wt% Cs/Al-PILC > 5wt% K/Al-PILC > 5wt% Na/Al-PILC > 5wt% Li/Al-PILC > Al-PILC > montmorillonite. The increased adsorption capacity of 5Cs/Al-PILC samples for CO2 was observed from 1.821 mmol/g before modification to 2.325 mmol/g, and the samples doped with alkali metals exhibited an excellent cycling performance (Fig. 10).
To further improve the capture efficiency of solid adsorbents for CO2, Sanz et al. [93] applied a dual-functional method that combines impregnation and grafting to modify silica to simultaneously obtain amino groups with high nitrogen insertion and high mobility. The results show that the CO2 adsorption rate of the bifunctionalized sample was increased by approximately 10%–20% compared with that of the impregnated sample. Therefore, some scholars have proposed the use of bifunctional methods for the modification of minerals [56–57,81]. Gómez-Pozuelo et al. [56] pregrafted 3-aminopropyl-trimethoxysilane or diethylenetriamine-propyl-trimethoxysilane on montmorillonite, sepiolite, bentonite, and other minerals and achieved bifunctional modification by PEI impregnation (Fig. 11). The results reveal that the CO2 adsorption capacity of bifunctional samples did not improve compared with that of samples prepared via a single impregnation or grafting method, but the value is close to that of nonfunctionalized montmorillonite. The main reason for this phenomenon is the likelihood of double-functionalized samples blocking the structural pores in minerals, which is not advantageous to CO2 diffusion, given the limitation of the porous structure of montmorillonite and sepiolite. To solve the clogging problem of adsorbents after dual functionalization, Cecilia et al. [57] treated acid-leached sepiolite by grafting it with APTES (AP) and dipping in PEI; they also proposed that the increased adsorption temperature benefited the rearrangement of the amine-rich polymer and CO2 diffusion; as a result, the amine sites were more fully in contact with CO2, thereby improving the CO2 adsorption capacity of the composite.
CO2 mineralization mimics the weathering process of Ca/Mg silicate minerals; this phenomenon was first proposed by Seifritz and can be expressed by the following general Eq. (1) [94]. Mineralization involves the use of natural minerals, including forsterite (Mg2SiO4), serpentine [Mg3Si2O(OH)4], wollastonite (CaSiO3), etc. [95–96]. These natural minerals are known for their abundant reserves, easy access, and stable products after mineralization [97]. Based on different reaction processes, CO2 mineralization can be divided into direct and indirect mineralization [20]. Direct mineralization refers to the mineralization of minerals via high-temperature and high-pressure operations under dry or wet conditions. Although the process is simple and uses fewer chemical reagents, it features high-temperature and high-pressure energy consumption, and directly obtaining high-value-added products is difficult, resulting in poor economic benefits. Therefore, some researchers have proposed indirect mineralization, in which minerals are first treated using acid, alkali, molten salt, ammonium salt, and other media to extract calcium and magnesium ions and then carbonated under certain conditions. Indirect mineralization has less harsh reaction conditions compraed with direct mineralization, and the obtained mineralized products are relatively pure and have certain economic value.
(Ca,Mg)xSiyOx+2y⋅zH2O(s)+xCO2(g)→x(Ca,Mg)CO3(s)+ySiO2(s)+zH2O(l/g) | (1) |
Based on different reaction modes between minerals and CO2, direct mineralization can be divided into dry and wet carbonation [44]. Dry carbonation refers to the direct reaction of CO2 with minerals at appropriate temperature and pressure. Wet carbonation involves the reaction of CO2 with minerals in an aqueous suspension [98], in which CO2 reacts with water to form bicarbonate and protons. Then, the protons dissolve metal ions, which can react with sodium bicarbonate to form carbonate precipitates in minerals.
Lackner et al. [95,97] first proposed the gas–solid carbonation of CO2. Through thermodynamic calculations, they showed that the gas–solid reaction of serpentine, olivine, and basalt with CO2 is an exothermic process. It can spontaneously react in nature, but the reaction rate is usually low. Wet carbonation was first proposed by O’Connor et al. [98]. Fig. 12 shows the process flow, and Eqs. (2)–(4) provide the reaction path; that is, CO2 dissolves in water to form carbonic acid. Then, the ore gradually dissolves and precipitates carbonate under the action of carbonic acid.
CO2+H2O→H2CO3→H++HCO−3 | (2) |
Mg2SiO4+4H+→2Mg2++H4SiO4(SiO2+2H2O) | (3) |
Mg2++HCO−3→MgCO3+H+ | (4) |
Huijgen et al. [99] reported the mechanism of direct mineralization of wollastonite in aqueous solutions. According to studies, throughout the carbonation process, factors such as reaction temperature, stirring speed, and liquid–solid ratio considerably influence the mineralization process of wollastonite. Reaction temperature affects the leaching of calcium and magnesium ions and nucleation efficiency of carbonates. At lower than optimal temperatures, the dissolution of calcium and magnesium ions on the outer surface during the mineralization process leaves a dense outer layer of SiO2. This condition affects the outward diffusion of calcium and magnesium ions from within the mineral and leads to leaching difficulties, whereas at high reaction temperatures, it influences the nucleation and growth of carbonates due to the lower activity of the generated carbonates. Stirring speed affects the leaching efficiency of calcium and magnesium ions, with the appropriate stirring speed being conducive to the carbonation of minerals. The liquid–solid ratio influences the thermal balance of the reaction process, which results in changes in the carbonation efficiency of minerals. At extremely low liquid–solid ratios, the viscosity of the reaction system increases, and problems such as difficult pumping and stirring occur, which drastically reduce the conversion efficiency. Meanwhile, an extremely high liquid–solid ratio leads to a reduced processing capacity of the reaction, which is not conducive to the industrialization of applications. In addition, improvement of the reaction rate of wet carbonation can be attained through pressure control in the reaction process, the addition of additives (such as NaHCO3 and NaCl [100–101]), ore treatment (such as heating activation [102–103], mechanical grinding [104–105]), and removal of the formed SiO2 inert surface layer. Table 5 shows the findings of the comparison of typical process parameters during the direct mineralization of CO2 by natural silicate minerals.
Mineral type | Particle size / μm |
Chemical additives | Reaction conditions | Mineralization rate / % |
Ref. | ||||
Liquid–solid ratio | Temperature / °C |
Pressure / MPa |
Time / h |
Stirring speed / (r·min−1) |
|||||
Wollastonite | 37–44 | 1 M NaHCO3 | 9 g/g | 110 | 8.6 | 6 | — | 90 | [100] |
Wollastonite | <38 | — | 5 kg/kg | 200 | 2.0 | 0.25 | 500 | 75 | [106] |
Olivine | <37 | 1 M NaCl, 0.5 M NaHCO3 | 5.6 g/g | 185 | 11.5 | 6 | — | 84 | [98] |
Olivine | <38 | 1 M NaCl, 0.64 M NaHCO3 | 2.33 g/g | 185 | 15 | 1 | — | 49.5 | [107] |
Olivine | <10 | 0.75 M NaCl, 0.5 M NaHCO3 | 5 mL/g | 190 | 10 | 4 | 1500 | 100 | [108] |
Olivine | <38 | 1 M NaCl, 0.64 M NaHCO3 | 9 mL/g | 185 | 6.5 | 6 | 1420 | 84.4 | [109] |
Serpentine | <37 | 1 M NaCl, 0.64 M NaHCO3 | 5.6 mL/g | 155 | 18.5 | 0.5 | 2000 | 78 | [98] |
Serpentine | <75 | 1 M NaCl, 0.64 M NaHCO3 | 2.33 g/g | 155 | 11.5 | 1 | — | 73.5 | [107] |
However, the entire process of direct mineralization still faces problems, such as harsh reaction conditions, high equipment requirements, high energy consumption during pretreatment, and low economic value of carbonization products. Therefore, the focus of the research on CO2 mineralization has gradually shifted to indirect mineralization.
In indirect mineralization, the active ingredients (such as calcium and magnesium ions) in mineral raw materials are first dissolved in the medium and then carbonated to form carbonates. Media screening is a key step in indirect mineralization. The reaction medium used must have two characteristics: (1) benefiting the leaching of calcium and magnesium ions; (2) easily recycled in the carbonation reaction process. The common media used in indirect mineralization mainly include hydrochloric acid, acetic acid, sodium hydroxide, molten salt, etc. Table 6 shows the findings of the comparison of typical process parameters in the indirect mineralization of CO2 by natural silicate minerals.
Mineral type | Particle size / μm | Chemical additives | Reaction conditions | Mineralization rate / % | Ref. | ||||
Liquid– solid ratio | Temperature / °C | Pressure / MPa | Time / h | Stirring speed / (r·min−1) | |||||
Serpentine | <500 | 2 M H2SO4, NaOH | 10 mL/g | 70 | 0.1 | 1 | 650 | 94 | [110] |
Lizardite | <150 | 2 M HCl, NH3 | 13.51 mL/g | 90 | 0.1 | 2 | 600–700 | 80 | [111] |
Serpentine | 75–150 | 1.4 M NH4HSO4, NH3 | 20 mL/g | 100 | 0.1 | 1 | 800 | 86 | [112] |
Serpentine | 125–212 | (NH4)2SO4 | — | 400 | 0.1 | 1 | — | 78 | [113] |
Wollastonite | — | 1.52 M CH3COOH, TBP | — | 80 | 4.0 | 1 | 500 | 50 | [114] |
Regarding reaction mediums, ores can be quickly dissolved in hydrochloric acid solution to form magnesium chloride, calcium chloride, and silica. This hydrochloric acid extraction process was developed in the 1940s and 1950s and was mainly used to extract magnesium from serpentine [115]. The reaction mechanisms are shown in Eqs. (5)–(8) [116]. The process mainly includes four steps: dissolution of serpentine by hydrochloric acid to form magnesium chloride, decomposition of magnesium chloride into Mg(OH)Cl at 150°C in water and excess hydrochloric acid, decomposition of Mg(OH)Cl into magnesium hydroxide and magnesium chloride, and formation of magnesium carbonate in CO2 atmosphere. The magnesium chloride and hydrochloric acid produced in this process can be recycled. Despite the effectiveness of this method, it still presents shortcomings, such as high energy consumption and strong equipment corrosion in the hydrochloric acid regeneration process, which limit its industrial application.
Mg3Si2O5(OH)4(s)+6HCl(aq)+H2O→3MgCl2⋅6H2O(aq)+2SiO2(s) | (5) |
MgCl2⋅6H2O(aq)→MgCl(OH)(aq)+HCl(aq)+5H2O | (6) |
2MgCl(OH)(aq)→Mg(OH)2(s)+MgCl2(aq/s) | (7) |
Mg(OH)2(s)+CO2(g)→MgCO3(s)+H2O | (8) |
In addition to hydrochloric acid extraction, other acids have been studied as mediators in the indirect mineralization process. Teir et al. [117] investigated the dissolution of natural serpentine in inorganic acids (HCl, H2SO4, HNO3, etc.), organic acids (HCOOH, CH3COOH, etc.), ammonium salt solutions ((NH4)2SO4, NH4Cl, NH4Cl, etc.), and alkali solutions (NaOH, KOH, NH3, etc.). The results reveal that all the tested acids can extract 3%–26% magnesium and 2%–16% iron from serpentine within 1 h at room temperature at 1–4 M acid concentration, with H2SO4 showing the highest extraction efficiency. Ammonium salt solution is the only selective solution used in magnesium extraction. However, its extraction rate is considerably lower than that of the acid used, and the alkaline solution fails to dissolve any measurable magnesium and iron. Teir et al. [110] used 2 mol/L H2SO4 to leach serpentine at 70°C. Almost all Mg can be extracted to the leaching solution, and the reaction equations are shown in Eqs. (9)–(11). In addition, approximately 65% of Fe and 3% of Si were extracted, and NaOH was used to remove Fe from the leaching solution for carbonation reaction. The carbonation rate of Mg can reach 94%. However, this process consumes extremely strong acids and alkalis, and the whole mineralization process is uneconomical.
Mg3Si2O5(OH)4(s)+3H2SO4(aq)→3MgSO4(aq)+2SiO2(s)+5H2O(l) | (9) |
FeO(s)+H2SO4(aq)→FeSO4(aq)+H2O(l) | (10) |
MgSO4(aq)+2NaOH(aq)+CO2(g)→MgCO3(s)+H2O(l)+Na2SO4(aq) | (11) |
O’Connor et al. [118] proposed that weak acids can be used instead of strong ones to lower energy consumption, and the recovery of media is relatively easy. The main reaction processes are shown in Eqs. (12)–(13).
CaSiO3(s)+2HAc(l)→Ca2+(aq)+2Ac−(aq)+H2O(l)+SiO2(s) | (12) |
Ca2+(aq)+2Ac−(aq)+CO2(g)+H2O(l)→CaCO3(s)+2HAc(l) | (13) |
Bao et al. [114] studied the carbon fixation process through the addition of tributyl phosphate (TBP) and magnesium acetate to acetic acid as the circulating medium (Fig. 13). The reaction processes are shown in Eqs. (14)–(19). This carbon-fixation process involves CO2 diffusion, dissolution, and ionization, calcium carbonate precipitation, and TBP acetic acid extraction, which can realize the recycling of acetic acid. However, the additives have the disadvantages of high energy consumption and complex processes during recovery, which increase the economic cost and additional carbon emissions. Despite the effectiveness of indirect mineral carbon sequestration via acid extraction, the cost is usually high. Thus, scholars should focus on exploring new extraction methods.
CO2(g)⟺CO2(aq) | (14) |
CO2(aq)+H2O(l)⟺H2CO3(aq) | (15) |
H2CO3(aq)→2H+(aq)+CO2−3(aq) | (16) |
Ca2+(aq)+CO2−3(aq)⟺CaCO3(s) | (17) |
Ac−(aq)+H+(aq)⟺HAc(aq) | (18) |
TBP(aq)+HAc(aq)⟺TBP⋅HAc(aq) | (19) |
Molten salt extraction is similar to HCl extraction, except for the use of molten salt (MgCl2·n(H2O)) instead of HCl as an extractant to reduce the energy consumption of CO2 during mineralization [119]. With serpentine as an example, the reaction processes are shown in Eqs. (20)–(23) [120]. Although the molten salt process consumes lower energy consumption than the HCl extraction process, rock impurities react with chloride ions, which causes the loss of chloride ions, and MgCl2 cannot be recovered. The complex process is difficult to operate and devoid of industrial application value.
3MgCl2⋅3.5H2O(l)+Mg3Si2O5(OH)4(s)→6Mg(OH)Cl(l)+9.5H2O(l)+2SiO2(aq) | (20) |
2Mg(OH)Cl(l)+nH2O(l)→MgCl2⋅nH2O(l)+Mg(OH)2(s) | (21) |
MgCl2⋅nH2O(l)→MgCl2⋅3.5H2O(l)+(n−3.5)H2O(l) | (22) |
Mg(OH)2(s)+CO2(g)→MgCO3(s)+H2O(g) | (23) |
Blencoe et al. [121] first proposed the use of NaOH in the extraction of Ca and Mg from metal silicate minerals to achieve mineral carbonation. With wollastonite as an example, the main reaction processes are shown in Eqs. (24)–(26). This extraction method has many problems, including (1) large energy consumption due to the long reaction time and high reaction temperature; (2) consumption of a large amount of alkali due to the high silicon content of the raw material; (3) difficulty of obtaining pure CaCO3 because of challenging product separation. Thus, NaOH is not an ideal medium for mineral carbonation in indirect mineralization carbonation technology.
CaSiO3(s)+NaOH(aq)→NaCaSiO3(OH)(s) | (24) |
2NaOH(aq)+CO2(aq)→Na2CO3(aq)+H2O | (25) |
Na2CO3(aq)+3NaCaSiO3(OH)(s)+H2O→4NaOH(aq)+CaCO3(s)+NaCa2Si3O8(OH)(s) | (26) |
The dissolution and carbonation of minerals usually consume large amounts of acid, alkali, and other media, which are difficult to recycle during use. To overcome the above limitations, Wang and Maroto-Valer [112] used a series of ammonium salt solutions (such as (NH4)2SO4, NH4Cl, and NH4HSO4) to dissolve serpentine; NH4HSO4 showed the highest effectivity in the extraction from serpentine samples; a comprehensive capture-and-storage process based on mineral carbonation using recyclable ammonium salts was also proposed (Fig. 14). The process initially used NH4HSO4 to dissolve magnesium ions from serpentine and then adjusted the pH through the addition of ammonia for the removal of impurities in the solvent. Second, after the removal of impurities, the solution was reacted with the NH4HCO3 formed by NH3 by capturing CO2 to form the precipitated carbonate. Finally, the solution mainly contained ammonium sulfate, which can be collected through evaporation and heated to regenerate ammonia. NH4HSO4 and NH3 can be recycled in this process. The results show that NH4HSO4 dissolved serpentine at 100°C for 3 h, and the extraction rate of Mg reached 100%. NH4HCO3 was reacted with Mg2+, and the resulting magnesium carbonate precipitation can reach up to 96%. However, in this process, ammonium sulfate consumed a notably high thermal decomposition energy, and the heat generated by mineralization reaction and mineral dissolution was difficult to utilize, resulting in energy loss.
To fully utilize the heat of acidolysis and mineralization, the reseachers [113,122] combined the regeneration and extraction process of ammonium bisulfate and used the heat of acidolysis and mineralization to accelerate the dissolution of serpentine and thermal decomposition of ammonium sulfate (Fig. 15). For the further improvement of the economics of the mineralization process, in addition to reduced energy consumption, mineralized (e.g., calcium carbonate and magnesium carbonate) and high value-added products[113,123–124] (e.g., porous silica and metals such as Cu and Ni) can be recovered.
Excessive CO2 emissions cause global warming, which substantially influences social, economic, and human production and development. In this context, the technology of CO2 capture and storage, which is deemed a highly effective method to mitigate CO2 emissions, has caught the interest of numerous scholars. This paper reviewed the progress of the research on CO2 capture and mineralization by natural minerals, with a focus on various methods that enhance CO2 capture and mineralization by natural minerals. Despite the considerable number of research advances, CO2 capture and mineralization by natural minerals still face many difficulties and challenges.
In terms of CO2 capture by natural minerals, the capability of adsorbents to capture CO2 is affected by the type of minerals, pretreatment methods, modifiers and modification conditions, adsorption conditions (concentration and flow rate of CO2, adsorption temperature and pressure, water content, etc.), and other factors. Meanwhile, pretreatment and modification methods directly affect the performance of solid adsorbents. Therefore, in the selection of pretreatment and modification methods, the inherent properties of minerals must be thoroughly considered, and the potential synergy between different modification techniques must be assessed. In addition, future research on CO2 capture by natural minerals will focus on the exploration of composite materials with high specific surface area and numerous active sites. Adsorption conditions affect the capacity of adsorbents to capture CO2. In determining adsorption conditions, simulation of the state of mixed gas during actual production is required to evaluate the adsorption capacity of solid adsorbents, and suitable mineral-based solid adsorbents must be prepared for different gas environments.
In terms of CO2 mineralization by natural minerals, indirect mineralization is milder than direct mineralization and does not require high-concentration and high-pressure CO2 as raw materials. The acquired mineralized products have high purity and possess economic value. However, several issues, including additional energy consumption, large consumption of media, and difficulty in utilizing heat generated by mineralization reactions and mineral dissolution, remain to be addressed. Therefore, future research on CO2 mineralization by natural minerals should focus on the development of new mineralization technologies (such as the combination of CO2 adsorption and mineralization) to reduce energy consumption during indirect mineralization, recycling of media, and full usage of high-value-added products.
Given strict carbon emission and neutral targets, it is necessary to promote the industrialization of CO2 capture and mineralization. However, the current research on natural mineral CO2 capture and mineralization mainly focuses on the laboratory stage. Therefore, practical applications should be simulated at a later stage, and related research should be accelerated from the laboratory stage to the pilot scale-up stage. In addition, the economic cost analysis of CO2 capture and mineralization is important. The development and industrialization of CO2 capture and mineralization technology can be further promoted through the establishment of a cost assessment system related to CO2 capture and mineralization.
This work was financially supported by the Beijing Natural Science Foundation, China (No. 2242055).
The authors declare that they have no conflicts of interest.
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[2] | Xiaoying Qian, Hong Yang, Chunfeng Hu, Ying Zeng, Yuanding Huang, Xin Shang, Yangjie Wan, Bin Jiang, Qingguo Feng. Effect of potential difference between nano-Al2O3 whisker and Mg matrix on the dispersion of Mg composites [J]. International Journal of Minerals, Metallurgy and Materials, 2023, 30(1): 104-111. DOI: 10.1007/s12613-022-2550-0 |
[3] | Wantong Chen, Wenbo Yu, Pengcheng Zhang, Xufeng Pi, Chaosheng Ma, Guozheng Ma, Lin Zhang. Fabrication and performance of 3D co-continuous magnesium composites reinforced with Ti2AlNx MAX phase [J]. International Journal of Minerals, Metallurgy and Materials, 2022, 29(7): 1406-1412. DOI: 10.1007/s12613-022-2427-2 |
[4] | Sheng Liu, Qing Yuan, Yutong Sima, Chenxi Liu, Fang Han, Wenwei Qiao. Wear behavior of Zn–38Al–3.5Cu–1.2Mg/SiCp composite under different stabilization treatments [J]. International Journal of Minerals, Metallurgy and Materials, 2022, 29(6): 1270-1279. DOI: 10.1007/s12613-020-2217-7 |
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1. | Chenguang Qian, Zhicheng Wang, Yi Wen, et al. Highly efficient CO2 capture on acid-treated sepiolite impregnated with mixed polyethyleneimine and diethanolamine. Green and Smart Mining Engineering, 2024, 1(4): 363. DOI:10.1016/j.gsme.2024.11.002 |
2. | Wei Wang, Liming Wu, Lin Chang, et al. Functionality developments in montmorillonite nanosheet: Properties, preparation, and applications. Chemical Engineering Journal, 2024, 499: 156186. DOI:10.1016/j.cej.2024.156186 |
Support | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||
Temperature / °C | Gas atmosphere | Pressure / MPa | |||
Kaolinite | 0 | CO2 | 3 | 0.29 | [48] |
25 | 3 | 0.14 | [48] | ||
25 | 0.1 | 0.06 | [49] | ||
15 | 1.7 | 0.14 | [50] | ||
Halloysite | 0 | CO2 | 3 | 6.17 | [51] |
75 | 0.1 | 0.08 | [52] | ||
Bentonite | 25 | CO2 | 0.1 | 0.13 | [53] |
30 | 0.14 | [54] | |||
Montmorillonite | 25 | CO2 | 0.1 | 0.16 | [55] |
45 | 0.22 | [56] | |||
Palygorskite | 45 | CO2 | 0.1 | 0.27 | [56] |
25 | 0.40 | [57] | |||
Sepiolite | 45 | CO2 | 0.1 | 0.93 | [56] |
50 | 0.41 | [58] |
Supports | Pretreatment | Structure property | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||||
Specific surface area / (m2·g−1) | Pore volume / (cm3·g−1) | Temperature / °C | Gas atmosphere | Pressure / MPa | |||||
Kaolinite | Without pretreatment | 8 | 0.01 | — | — | — | — | [65] | |
6 M HCl, 4 h | 38 | 0.07 | |||||||
5 M NaOH, 4 h | 16 | 0.03 | |||||||
Halloysite | Without pretreatment | 63.4 | 0.24 | — | — | — | — | [66] | |
850°C, 4 h + 6 M HCl, 6 h | 366.4 | 0.55 | |||||||
Bentonite | Without pretreatment | 81 | 0.10 | 30 | CO2 | 0.1 | 0.14 | [54] | |
3 M HCl, 3 h | 227 | 0.31 | 0.34 | ||||||
Montmorillonite | Without pretreatment | 72 | 0.16 | — | — | — | — | [65] | |
6 M HCl, 4 h | 253 | 0.71 | |||||||
5 M NaOH, 4 h | 183 | 0.56 | |||||||
Sepiolite | Without pretreatment | 79 | 0.01 | 50 | CO2 | 0.1 | 0.41 | [58] | |
2 M HCl, 24 h | 182 | 0.21 | 0.65 | ||||||
Without pretreatment | 42.74 | 0.08 | 75 | CO2 | 0.1 | 0.27 | [64] | ||
2 M HCl, 6 h | 168.96 | 0.14 | 0.41 | ||||||
Without pretreatment | 42.74 | — | 70 | CO2 | 0.1 | 0.30 | [67] | ||
4 M HCl, 6 h | 320.31 | 0.50 |
Support | Pretreatment | Amine loading / wt% | Adsorption conditions | Regeneration conditions | Cycle number |
CO2 uptake / (mmol·g−1) | Loss / % |
Ref. | |||||
Temperature / °C |
Atmosphere | Pressure / MPa |
Temperature / °C |
Atmosphere | X | Xresidue | |||||||
Kaolinite | Intercalation | 6.35 (HXA) | 0 | CO2 | 3 | 105 | N2 | — | 0.75 | — | — | [48] | |
12.6 (DDA) | 0.50 | ||||||||||||
41.7 (ODA) | 0.28 | ||||||||||||
Without pretreatment | 20 (MEA) | 25 | CO2 | 0.1 | 100 | N2 | — | 2.94 | — | — | [49] | ||
33.3 (EDA) | 1.35 | ||||||||||||
50 (4MEA–1EDA) | 3.38 | ||||||||||||
Halloysite | Without pretreatment | 44.4 (Formamide) | 0 | CO2 | 3 | 250 | — | — | 6.1 | — | — | [51] | |
25 | 3.4 | ||||||||||||
Without pretreatment | 30 (PEI) | 75 | CO2 | 0.1 | 75 | N2 | 10 | 1.5 | 1.4 | 6.6 | [52] | ||
Hydrochloric acid treatment | 50 (PEI) | 85 | 40vol%N2 + 60vol%CO2 | 0.1 | 110 | N2 | 10 | 2.75 | 2.66 | 3.2 | [66] | ||
Hydrochloric acid treatment | 30 (PEI) | 25 | Dry air | 0.1 | 80 | — | 50 | 1.24 | 1.20 | — | [77] | ||
Palygorskite | Hydration treatment | 37 (PEI) | 45 | CO2 | 0.1 | 110 | Ar | — | 1.52 | — | — | [56] | |
Sepiolite | Hydrochloric acid treatment | 47 (PEHA) | 50 | CO2 | 0.1 | 105 | N2 | 10 | 2.47 | 2.23 | 9.8 | [58] | |
15vol%CO2 + 85vol%N2 (water content 5vol%) | — | 2.94 | — | — | |||||||||
Hydrochloric acid treatment | 50 (PEI) | 75 | 40vol%N2 + 60vol%CO2 | 0.1 | 110 | N2 | 10 | 2.48 | 2.31 | 6.8 | [64] | ||
Hydrochloric acid treatment | 60 (TEPA) | 60 | 1vol%CO2 + 99vol%N2 (water content 1vol%) | 0.1 | 90 | N2 | 10 | 3.8 | 3.5 | 7.9 | [78] | ||
Bentonite | Without pretreatment | 30 (PEI) | 75 | CO2 | 0.1 | 75 | He | 5 | 1.06 | 1.02 | 3.7 | [53] | |
Sulfuric acid treatment | 50 (TEPA) | 75 | CO2 | 0.1 | 100 | N2 | 10 | 2.54 | 2.43 | 4.3 | [79] | ||
15vol%CO2 + 4.5vol%O2 + 80.5vol%N2 (water content 18vol%) | — | 4.31 | — | — | |||||||||
Montmorillonite | Hydration treatment | 37 (PEI) | 45 | CO2 | 0.1 | 110 | Ar | — | 1.03 | — | — | [56] | |
Pillaring and calcination | 60 (TEPA) | 25 | CO2 | 0.1 | — | — | — | 1.64 | — | — | [80] | ||
60 (PEI) | 1.45 | ||||||||||||
Hydrochloric acid treatment | 50 (PEI) | 75 | CO2 | 0.1 | 100 | N2 | 10 | 2.54 | 2.43 | 4.4 | [65] | ||
15vol%CO2 + 4.5vol%O2 + 80.5vol%N2 (water content 3vol%) | — | 3.23 | — | — | |||||||||
Without pretreatment | 20–30 (ODA) | 25 | CO2 | 5 | — | — | — | 7.16 | — | — | [81] | ||
Note: X—CO2 uptake for fresh sample; Xresidue—CO2 uptake for cycle sample. |
Support | Pretreatment | Modifier | Adsorption conditions | CO2 uptake / (mmol·g−1) | Ref. | ||
Temperature /°C | Atmosphere | Pressure / MPa | |||||
Halloysite | Microwave-assisted acid treatment | APTES | 25 | CO2 | 0.1 | 0.11 | [52] |
Palygorskite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.86 | [56] |
DT | 1.29 | ||||||
Microwave-assisted acid treatment | APTES | 25 | CO2 | 0.1 | 0.75 | [57] | |
Calcination and hydrochloric acid treatment | APTES | 40 | Air | 0.1 | 1.81 | [84] | |
Sepiolite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.99 | [56] |
DT | 1.39 | ||||||
Microwave-assisted acid | APTES | 25 | CO2 | 0.1 | 1.0 | [57] | |
Montmorillonite | Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.75 | [56] |
DT | 1.13 | ||||||
Inserted layer treatment | AEAPTS | 95 | 15vol%CO2 +85vol%N2 | 0.1 | 1.80 | [86] | |
Bentonite | Hydrochloric acid treatment | APTES | 30 | CO2 | 0.1 | 0.66 | [54] |
Hydration treatment | APTES | 45 | CO2 | 0.1 | 0.98 | [56] | |
DT | 0.73 | ||||||
NaCl and CTAB treatment | APTES | 25 | CO2 | 0.1 | 1.14 | [80] | |
Note: DT represents N1-(3-trimethoxysilylpropyl) diethylenetriamine. |
Mineral type | Particle size / μm |
Chemical additives | Reaction conditions | Mineralization rate / % |
Ref. | ||||
Liquid–solid ratio | Temperature / °C |
Pressure / MPa |
Time / h |
Stirring speed / (r·min−1) |
|||||
Wollastonite | 37–44 | 1 M NaHCO3 | 9 g/g | 110 | 8.6 | 6 | — | 90 | [100] |
Wollastonite | <38 | — | 5 kg/kg | 200 | 2.0 | 0.25 | 500 | 75 | [106] |
Olivine | <37 | 1 M NaCl, 0.5 M NaHCO3 | 5.6 g/g | 185 | 11.5 | 6 | — | 84 | [98] |
Olivine | <38 | 1 M NaCl, 0.64 M NaHCO3 | 2.33 g/g | 185 | 15 | 1 | — | 49.5 | [107] |
Olivine | <10 | 0.75 M NaCl, 0.5 M NaHCO3 | 5 mL/g | 190 | 10 | 4 | 1500 | 100 | [108] |
Olivine | <38 | 1 M NaCl, 0.64 M NaHCO3 | 9 mL/g | 185 | 6.5 | 6 | 1420 | 84.4 | [109] |
Serpentine | <37 | 1 M NaCl, 0.64 M NaHCO3 | 5.6 mL/g | 155 | 18.5 | 0.5 | 2000 | 78 | [98] |
Serpentine | <75 | 1 M NaCl, 0.64 M NaHCO3 | 2.33 g/g | 155 | 11.5 | 1 | — | 73.5 | [107] |
Mineral type | Particle size / μm | Chemical additives | Reaction conditions | Mineralization rate / % | Ref. | ||||
Liquid– solid ratio | Temperature / °C | Pressure / MPa | Time / h | Stirring speed / (r·min−1) | |||||
Serpentine | <500 | 2 M H2SO4, NaOH | 10 mL/g | 70 | 0.1 | 1 | 650 | 94 | [110] |
Lizardite | <150 | 2 M HCl, NH3 | 13.51 mL/g | 90 | 0.1 | 2 | 600–700 | 80 | [111] |
Serpentine | 75–150 | 1.4 M NH4HSO4, NH3 | 20 mL/g | 100 | 0.1 | 1 | 800 | 86 | [112] |
Serpentine | 125–212 | (NH4)2SO4 | — | 400 | 0.1 | 1 | — | 78 | [113] |
Wollastonite | — | 1.52 M CH3COOH, TBP | — | 80 | 4.0 | 1 | 500 | 50 | [114] |