A review on removal CO₂, SO₂, and H₂S from flue gases using zeolite based adsorbents

Abstract

This review was carried out on removal of flue gases (CO₂, SO₂, and H₂S) that are emitted from various sources. Burning solid fuels for heat, such as natural gas, gasoline, and coal or biomass, results in the production of flue gas. Adsorption of flue gases by utilizing Zeolites was properly explained, including the zeolite synthesis technique, characteristics, zeolite sensitivity, variables that influence the adsorption process, zeolite efficiency, and cost-effectiveness. Zeolites have a notable adsorption capability for CO₂, SO₂, and H₂S, despite their major disadvantage of being poor long-term durability and stability. Remarkable advancements are being made in present-day CO₂ capture methods, especially concerning the use of zeolites. Zeolites are unique materials with great potential for CO₂ collection because of their tiny holes, high porosity, structural variety, and recyclability. Since atmospheric CO₂ concentrations are already more than 415 parts per million, it is imperative to limit future releases of this gas and keep it out of the carbon cycle. The elimination of SO₂ by the majority of zeolites was shown to rise with temperatures up to 200 °C and subsequently decrease at higher temperatures. It was also shown that SO₂ uptake and breakthrough time are significantly affected by drying zeolite using microwave and traditional heating methods. The sorption capacity and sorbent regenerability of SO₂ are sensitive to the presence of other gases in the flue gas, such as O₂, CO₂, NOx, and water vapor, as well as the reaction temperature. Zeolites must possess strong sulfur loading capacity, good regenerability, and a stable structure to be an effective adsorbent for the removal of H₂S. Regarding the utilization of zeolites as adsorbents for the flue gases, new developments, and continuing difficulties have been concluded in this review work.

Article Highlights

• Efficacy of Zeolites for flue gas removal.

• Several challenges in Zeolite utilization.

• Further research on Zeolites' performance.

1 Introduction

1.1 Overview of flue gases and air pollution

Many companies, including those in food manufacturing, paper and pulp, pharmaceuticals, fuel extraction, and main metal industries, use boilers to generate steam [1]. Boiler requires heat from combustible burning. To successfully use fuel and reduce furnace losses, it is necessary to understand the energy balance or energy economy of the boiler [2]. Flue gases from heating fuel that do not convert to water or steam are the boiler's main source of waste. The temperature of the exhaust gas exiting a furnace varies from 150 to 250 ºC, or roughly 9 to 30% of the boiler's total energy production. Dincer et al. [3] stated that energy analysis, which is founded on the first rule of thermodynamics, is a common technique for determining how much energy is used in practical processes. The quantity and rate of energy or resource deterioration during a process, however, cannot be determined by energy analysis [3].

The collection of thermal power plants that use coal as their main energy source to produce electricity includes coal-fired power plants. Both dark coal and hard coal are used as fuel. In a mill, the coal is broken up, desiccated, and pulverized. The pulverized coal is fired into the furnace along with prepared combustion air and consumed at temperatures of about 1000 ºC [4]. A boiler's combustion of fuel can produce a variety of exhaust gases. Their formation is typically influenced by the fuel's makeup and the burning process. Temperature, agitation, and duration are three elements that have a significant impact on the burning process. Each factor's significance depends on the boiler's construction and the material being consumed [5]. Even though the heat generated through combustion is helpful for a broad range of uses, air pollutants, which are also made as a byproduct of the combustion process, must be regulated. CO₂,SO₂, NO_x, H₂S, and CO are typical atmospheric pollutants [6].

One significant source of carbon dioxide is the flue gas that power plants release. Flue gas is a blend of chemicals produced by burning and other processes. 33 to 40% of all world carbon dioxide emissions are caused by these pollutants. A predominant constituent of exhaust gas, exceeding 70%, is nitrogen, with the remaining portion comprising carbon dioxide, recognized as a pollutant [7]. The paramount contributor to global warming is carbon dioxide, the quintessential greenhouse gas. The swift escalation in atmospheric carbon dioxide concentrations is a consequence of the combustion of carbon-based sources for power generation. Given the alarming surge in global temperatures, urgent strides in technological innovation are imperative to curtail carbon dioxide emissions [8].

The proliferation of human activities and economic expansion has triggered a substantial surge in the consumption of fossil fuels, resulting in an annual increase in atmospheric carbon dioxide levels. Data from the UK Met-Office revealed a new record high for atmospheric CO₂ levels, underscoring the critical need for proactive measures to address this escalating environmental challenge [9]. Fuels containing sulfur are burned, producing SO₂ as a byproduct. The two main industrial sources of sulfur pollution are the burning of fossil fuels and the processing of metallic materials. In the environment, sulfur dioxide oxidizes and becomes sulfuric acid. Anthracite consumed in coal-fired power stations had an average sulfur concentration of 2.3% [10]. Utilizing black with less sulfur decreased this figure to 1.34%. Although switching to low-sulfur fuels can decrease SO₂ pollution and allow utilities to abide by regulatory SO₂ emission limits, it may harm furnace efficiency. If moving to low-sulfur coal is ineffective in bringing SO₂pollution down to a manageable level [11].

Even at low concentrations, hydrogen sulfide (H₂S) is a gas that is combustible, corrosive, and deadly. Therefore, to prevent and mitigate the harmful effects of H₂S on humans (which can lead to respiratory failure and asphyxiation) and other facilities (which can cause industrial equipment and pipeline corrosion), it must be captured and removed from a variety of emitting sources, including gas processing facilities, natural emissions, sewage treatment plants, landfills, and other industrial plants [12].

At the moment, eliminating species that contain sulfur is a strategic concern [13, 14] based on environmental laws and efforts to find new, clean energy sources like fuel cells [15]. Gaseous fuel must have almost no mercaptans or sulfides to be used in fuel cells. The reforming catalysts are poisoned by those species, which reduces cell longevity and raises energy expenses. Significant efforts are undertaken to filter air from locations where anaerobic digestion might occur, such as from municipal wastewater treatment plants, because the sulfur-containing gases are hazardous and have a very low odor threshold [16,17,18]. By adding a calcium-based sorbent, such as dolomite or limestone, to the bed of a fluidized-bed gasifier, in-bed sulfur capture can be achieved. The sulfur is then removed from the system together with the bottom ash in the form of CaS. The pressurized H₂S uptake by a variety of chemically and physically distinct limestones and dolomites, usually for those in a pressurized fluidized-bed gasifier operating at 2 MPa at 950 °C. Two pCO₂ values were tested using a pressurized thermobalance. This allowed for the analysis of the sulfidation of both calcined and uncalcined sorbents. For half-calcined dolomites and uncalcined limestones, the effect of pH₂S was also studied [19].

There are many review papers on removal of CO₂, SO₂ and, H₂S. However, those papers mainly focus only on the capturing mechanism. On the other hand, this review paper discusses in detail not only the removal mechanism (adsorption–desorption and adsorption kinetics) but also the efficiency, cost-effectiveness of the flue gases. In addition to that, readers could be able get the method of capturing/removal of these flue gases in a single paper and compare and contrast the removal efficacy of zeolites. To that end, this paper paves the way to industries, policy makers, engineers, entrepreneurs, businessmen to give a detail knowledge about the adsorption of these types of flue gases using zeolites in practical applications (large scale) and consequently will contribute in air pollution reduction and controlling of global warming. The paper concludes by highlighting both the advancements and continuing difficulties in utilizing zeolites as adsorbents for flue gas removal. This identification of challenges opens avenues for future research and development in this field, contributing to ongoing efforts to address environmental pollution and climate change.

1.2 Defining zeolites: exploring types, structures, properties, and applications

Zeolites, intricate and highly structured hydrated aluminosilicate crystals, manifest a distinct 3-dimensional TO₄ tetrahedral architecture within alkalies and alkaline earth. A notable characteristic of their framework is the generation of a negative charge facilitated by the isomorphous non-equivalent replacement of Al³⁺ for Si⁴⁺ in the tetrahedra [20]. This structural arrangement is further nuanced by the connectivity of oxygen ions, which link two cations and are simultaneously shared by two tetrahedrons [21].

The complexball of isomorphic substitution becomes pivotal, especially as some Si⁴⁺ ions are replaced by Al³⁺, giving rise to a net negative charge in the tectosilicate structure. This charge asymmetry is attributed to the distinct formal valence of (AlO₄)⁵⁻ and (SiO₄)⁴⁻ tetrahedrons, typically positioned on an oxygen anion linked to an aluminum cation [22]. To restore equilibrium, cations, predominantly Na⁺, K⁺, Ca²⁺, and Mg²⁺, inhabit the pathways within the zeolite structure, countering the prevailing negative charge.

The delicate balance within the channels allows for the tenuous retention of exchangeable cations, swiftly replaced by environmental cations, as indicated by Osacký et al. [23]. Consequently, counter-ions are strategically positioned on the exterior surface of zeolites, forming associations through weak electrostatic interactions to neutralize the inherent negative charges in the structure [23]. These counter-ions, marked by interchangeability, populate the spaces formed by the 3-dimensional link tetrahedral framework, contributing to the diverse array of polyhedral structures that shape the expanded frameworks defining various zeolite crystal patterns [24].

The composition of zeolites can be expressed through the general formula Mx/n [(AlO₂) x(SiO₂) y]. zH₂O, where n represents the charge of the metal cation M. The specific values of x, y, and z vary based on the type of zeolite under consideration. For instance, zeolite NaA, also known as LTA (Linde Type A), is represented by the formula Na₁₂ [(AlO₂)₁₂(SiO₂)₁₂]0.27H₂O. This indicates that zeolite LTA consists of 12 tetrahedra within each cell unit, housing 12 Na atoms and 27 H₂O molecules [25].

Al-dahri et al. [26] produced Linde-type A zeolite (LTA) from coal fly ash (CFA) and employed it as a sorbent to extract acidic dye (Acid red 66, AR66) from an aqueous solution. Through batch experiments, they assessed the effectiveness of the synthesized LTA in adsorbing AR66 from the solution [26]. Linde-type A (LTA) and Na-P type (ZP) zeolites were synthesized from the same raw material, namely coal fly ash, which is released from coal power plants, by Al-Dahri et al. [27] using microwave irradiation. The prepared zeolites underwent characterization and modification using a cationic surfactant for the removal of methyl orange dye from its aqueous solution [27]. Zeolites X and Y fall into the faujasite (FAU) group, encompassing faujasite-Na, faujasite-Mg, or faujasite-Ca. These zeolites share a common basic formula: (Na, Ca, Mg)_3.5 [A_l7Si₁₇O₄₈]0.32(H₂O). The structures of these three zeolite types are illustrated in Fig. 1 [25].

Fig. 1 

Structures and micropore networks of four distinct zeolites (Examples of building units and pore/cage sizes of three zeolite structures: zeolite A (LTA), sodalite (SOD), and faujasite (FAU)—zeolites X, Y) [25]

Porosity in zeolites is a consequence of a complex network of interconnected cages that extend to the particle's exterior surface, typically exhibiting dimensions less than 2 nm, thereby classifying zeolites as microporous materials [21]. This microporous nature stems from the solid and uniform openings within zeolite structures, dictated by the precise arrangement of lattice elements forming their walls. The inherent selectivity of zeolites arises from this uniformity, allowing only molecules with specific dimensions and shapes to be abundantly sorbed. This characteristic has led to the designation of zeolites as molecular sieves, a concept articulated by Barrer [28]. The architectural blueprint of zeolite structures is predominantly shaped by the repetition of secondary/auxiliary construction units (SBUs), and as of the present, there exist 19 distinct SBUs [28]. Another insightful approach to classifying zeolites involves the consideration of their pore apertures and channel dimensions. This classification scheme discerns between small pore zeolites, characterized by 8-membered-ring pores, medium pore zeolites with ten-membered rings, and large pore zeolites featuring 12-membered-ring pores. A recent addition to this categorization includes the extra-large pore zeolite group [29, 30]. This categorization not only aids in understanding the structural diversity of zeolites but also facilitates the comparative analysis of their adsorptive, molecular screening, and catalytic characteristics.

The expansiveness of channels within zeolite frameworks is intricately linked to the size of the cations they can accommodate, with larger cations correlating to broader channels, as elucidated by Osacký et al. [31]. The inherent versatility of zeolites, stemming from their adjustable framework composition, particle size, and channel diameter, renders them highly desirable for a myriad of applications, encompassing adsorption, ion exchange, removal of anionic dyes, and catalysis [26, 31].

The utility of zeolites spans a wide spectrum in the chemistry industry, with applications ranging from water filtration to gas separation, and from hydrocarbon processing to ion adsorption and exchange. The ability to tailor the framework composition allows for customization based on specific needs, making zeolites indispensable in diverse chemical processes. The adjustable particle size and form further contribute to their adaptability in various industrial contexts. In essence, zeolites serve as versatile tools in the chemical industry, providing solutions for an array of challenges. Their ability to selectively adsorb, exchange ions, and catalyze reactions underscores their importance in applications spanning from environmental remediation to the synthesis of valuable chemicals [32, 33]. The interplay between framework composition and channel dimensions not only enriches our understanding of zeolite behavior but also enhances the precision with which these materials can be harnessed for specific chemical processes.

2 Flue gas treatment

Power plants and various industrial facilities face the imperative of complying with stringent clean-air regulations imposed by national authorities. To mitigate the release of pollutants into the atmosphere, these sites are mandated to employ exhaust gas treatment methods. Such methods, exemplified by the use of equipment like scrubbers and electrostatic precipitators, have demonstrated efficacy in reducing the emission of contaminants by 90% or more. However, the implementation and operation of these pollution control technologies can entail substantial costs, making them a significant financial burden for power plants and industrial installations. Moreover, the regulatory landscape surrounding exhaust gas cleaning often becomes a focal point for prolonged legal disputes. The complexity and variability of regulations further add to the challenges faced by these facilities, with different plants necessitating specific treatments and certain nations imposing more stringent regulations than others [34].

Among the various procedures employed for pollutant removal, sulfur dioxide elimination stands out as a critical concern. The majority of these procedures involve the use of scrubbers, and wet scrubbers are the preferred choice for many American sites. Wet scrubbers utilize a combination of alkaline sorbent, typically composed of limestone or lime, and saltwater to effectively clear pollutants from the exhaust gases [35]. This underscores the intricate and site-specific nature of exhaust gas treatment, reflecting the multifaceted challenges and considerations faced by industries striving to balance regulatory compliance with operational efficiency.

3 Mechanism of adsorption of flue gases

Adsorption mechanisms can be categorized into two types based on the nature of interactions: physisorption and chemisorption. Physisorption refers to the physical adsorption of molecules, which occurs due to physical forces such as dipole–dipole interactions, electrostatic attractions, polar interactions, hydrophobic associations, or Van der Waals forces. The bond energy in physisorption ranges from 8 to 41 kcal/mol. On the other hand, chemisorption involves the chemisorbed attachment of molecules through chemical bonds, including covalent bonds, ionic bonds, or metallic bonds. The bond energy in chemisorption is approximately 60 to 418 kcal/mol [36]. According to reports, the zeolite adsorptive desulfurization of H₂S relies on either physisorption or chemisorption processes. Two chemisorption processes for H₂S were identified: sulfide-sorbent π-complexation and sulfur-metal (S-M) direct coordinate bonding. Zeolites can undergo ion exchange with different metal ions such as K⁺, Ag⁺, Cu⁺, Zn²⁺, and Ni²⁺,resulting in the formation of π-complexation between the metal ions and sulfur [37].

Sigot et al. [38] discovered a technique for the adsorptive removal of H₂S using zeolite 13X. This mechanism encompasses both physisorption and chemisorption processes, as outlined below [38].

• i. Adsorption at the surface of zeolite: (H₂S(g) → H₂S (ads)), physisorption.

• ii. Dissolution: (H₂S(ads) → H₂S (aq))

• iii. Dissociation reaction: (H₂S (aq) + H₂O (l) → HS⁻ (aq) + H₃O⁺ (aq)), due to the presence of water within the pores of the zeolite.

• iv. Oxidation reaction: (HS⁻(aq) + O (ads) → S (ads) + OH⁻ (aq)), due to the presence of adsorbed oxygen inside the pores of zeolites.

• v. Formation sulfur polymers: xS _(ads) → Sx _(ads).

Tran et al. [39] demonstrated that in addition to physisorption, the adsorption of H₂S on cobalt (II)-exchanged NaX zeolite takes place through an acid–base process (chemisorption), as described by Eq. 1 [39]:

$${\text{H}}_{{2}} {\text{S}} + \left[ {{\text{Co}}\left( {{\text{O}}_{{\text{x}}} } \right){6}} \right]^{{{2} + }} {\text{X}}^{{{2} - }} {\text{CoS}} \downarrow + {\text{ H}}_{{2}} {\text{X}} + {\text{O}}_{{\text{x}}}$$

(1)

The chemisorption process of H₂S by the ion-exchanged ZnX zeolite was proposed by Long et al. [41]and is demonstrated in reactions 1 (coordinate bonding) and 2 [40]:

• 1. Zn²⁺  + H₂S → (Zn-SH₂)²⁺ads

• 2. (Zn-SH₂)²⁺ads → ZnS + 2H⁺

The process of reactive H₂S adsorption by Cu-ETS-2 (Engelhard Titanosilicate), a molecular sieve zeolite, was studied by Yazdanbakhsh et al. [41]. The suggested adsorption reaction (chemisorption) was provided by Eqs. 2 and 3, respectively, for the range of adsorption temperatures between 250 and 650 °C and above 650 °C [41]:

$${\text{CuO}} + {\text{H}}_{{2}} {\text{S}} \to {\text{CuS}} + {\text{H}}_{{2}} {\text{O}}$$

(2)

$${\text{2CuO}} + {\text{H}}_{{2}} {\text{S}} + {\text{H}}_{{2}} \to {\text{Cu}}_{{2}} {\text{S}} + {\text{2H}}_{{2}} {\text{O}}$$

(3)

Using IR, UV–vis, and NMR spectroscopies, the mechanisms of H₂S adsorption on FAU and Linde-Type-A (LTA) zeolites were investigated [42,43,44,45]. Two different adsorption mechanisms, chemisorption, and physisorption, were identified in these investigations. Equation 4 illustrates the initial mechanism (step 1, chemisorption), which is thought to be the dissociation of the H₂S molecule on FAU surfaces [23]. Low pressures were used, and the initial H₂S molecules that were adsorbed were dissociated [44].

$${\text{H}}_{{2}} {\text{S2H}}^{ + } + {\text{S}}^{{{2} - }}$$

(4)

The equation indicates that more H₂S molecules adsorbed on ion-exchanged NaX zeolite underwent dissociation as the pressure was increased. It was shown that this mechanism is dependent on the Si/Al ratio and the H₂S pressure, meaning that it works only for Si/Al ratios ≤ 2.5 and at pressures as high as 6.66 mbar [43].

$${\text{H}}_{{2}} {\text{S}} + {\text{Na}}^{ + } {\text{Na}}^{ + } {\text{HS}}^{ - + } + {\text{H}}^{ + }$$

(5)

H₂S molecules adsorb on weaker zeolite surface sites as the amount of adsorbent material covered rises. Thus, the physisorption mechanism involved the coordination of H₂S molecules to surface Na⁺ cations and the formation of an H-bond with the zeolite framework's oxygen atom (Eq. 5) [43, 45].

3.1 Adsorption mechanism of CO₂

Physisorption is the primary mechanism of CO₂ adsorption on zeolite. Studies indicate that CO₂ physisorbs on alkali zeolites in a linear orientation through ion–dipole interaction, as structure A illustrates [46, 47].

$$\left( {\text{metal ion}} \right)^{{{\text{x}} + }} \ldots \ldots \ldots \ldots \ldots \ldots^{{{\updelta } - }} {\text{O = C = O}}^{{{\updelta } + }} \quad ({\text{A}})$$

These zeolites also included CO₂ that was more firmly bonded, in addition to physical adsorption. As can be seen structures B and C, these bent adsorbed CO₂ sites are connected to bi-coordination. Structure B depicts a carboxylate species that forms on zeolite as a result of polarized CO₂ molecules interacting with a lattice O ion and a less protected cation. Eventually, structure B changed into structure C, which is a carbonate structure. On the other hand, a particular kind of Na⁺ ion and an oxygen atom in the neighboring lattice may be necessary for the formation of a carbonate structure. To polarize CO₂ molecules, a less-shielded cation with a high enough electron potential and pore diameter is needed. Carbonate species were formed as a result of interaction with a charged lattice O ion. On the other hand, in zeolites with smaller pore sizes, such as 4A and 5A, these sites might not be readily accessible if type III active sites are present inside the pores for zeolite type X [46, 47].

3.2 Adsorption mechanism of SO₂

It has been noted that SO₂ may be physically bound (physisorption) in the zeolite cavities, a process that is readily regenerable. The infrared (IR) analysis of SO₂ adsorption on HY zeolites revealed an S–O stretching vibration band around 1320–1340 cm^–1, along with a 20–40 cm^–1 shift in frequency, indicating physical adsorption of SO₂. Because SO₂ molecules couldn't enter the tiny HY spaces, physisorption took place [48]. The ion–dipole interaction of Na⁺ with SO₂ caused SO₂ to be chemically adsorbed (chemisorption) on NaY zeolite, as demonstrated by IR and TPD experiments. The amount of SO₂ adsorbed chemically increased as the Si/Al ratio dropped. However, the amount of SO₂ physically adsorbed on zeolite rose as the Si/Al ratio increased [49]. The study examined the adsorption of SO₂ on NaY zeolite at a temperature range of 25 to 200 °C. An analysis using infrared (IR) spectroscopy verified that SO₂ molecules were adsorbed chemically via hydrogen bonds. This link was created by engaging with NaY's accessible OH groups, as seen in Fig. 2 [50].

Fig. 2

An example of the SO₂ adsorption mechanism on zeolite Y [50]

The formation of Lewis basic sites in zeolite structure can be attributed to the oxygen atom bridging in the Si–O–Al moiety. Through donor–acceptor contact, this Lewis basic oxygen can form a bond with the sulfur atom of an SO₂ molecule [51,52,53]. According to Nasluzov et al. [54] SO₂ and the HY framework may interact in two different ways. These are donor–acceptor interactions and hydrogen bonding. Figure 3 depicts potential SO₂ chemisorption structures with Y zeolite [54].

Fig. 3

Model structures of (a) SO₂ adsorbed hydrogen bonding; (b) Donor–acceptor interaction causes SO₂ to be adsorbed, and (c) SO₂ interacts with the zeolite's bridging hydroxyl and basic oxygen center [54]

The adsorption of SO₂ on NaY zeolite, which has a silicon/aluminum ratio of 2.6, was investigated by Kirschhock et al. [55]. The author claimed that hydrogen bonding and donor–acceptor interactions caused sulfite to develop and stabilize on zeolite [55].

4 Flue gas treatment using zeolites as adsorbent

Growing attention has been given to the issue of energy and environmental sustainability sectors, particularly to the control of burning exhaust gas emissions from coal-fired power plants, industrial furnaces, and kilns, as a result of the fast societal and economic growth [56,57,58]. After urbanization, pollution levels reached a high, which in turn had an impact on human, vegetation, and animal survival. The typical climate change of the world, the weather, water level, animal areas, and species are all negatively impacted by environmental contamination [59,60,61,62]. Since its content exceeded 400 parts per million, carbon dioxide emissions' effect on climate change has gotten stronger every year [63]. This illustrates the important role that human actions play in the pollution of the environment. Carbon capture and storing, also known as sequestering, is among the most efficient ways to fight against global climate change among the available strategies [64]. Carbon dioxide is widely absorbed by using carbonaceous materials like activated carbon [65], carbon nanotubes [66], and graphene [67], and also other solid adsorbents like mesoporous silica [68], zeolites [69], and metal–organic frameworks [70]. Due to their excellent thermo-chemical stability, responsive method of production, cost-effectiveness, and adaptable physical characteristics, zeolite-based compounds have among these adsorbents displayed extraordinary exhaust gas capture properties. A summary of CO₂ removal technologies has been shown in Table 1.

Table 1 Summary of CO₂ capturing methods: benefits and drawbacks

Various researchers have delved into the study of eliminating or separating flue gases utilizing zeolites emitted from diverse sources. Scholars across different disciplines have dedicated their efforts to exploring the potential of zeolites in facilitating the removal or separation of these flue gases. The utilization of zeolites, crystalline aluminosilicate minerals with well-defined porous structures, presents a promising avenue for addressing the challenges associated with flue gas emissions. Researchers have examined the unique adsorption and ion-exchange properties of zeolites to develop innovative methods for capturing and separating specific components within flue gases.

4.1 Utilizing various zeolites for adsorption in carbon dioxide removal

The existence of humans is threatened by high CO₂ releases, which cause global climate change and extreme weather [78]. The majority of human energy-related carbon dioxide emissions, or up to 44% of total emissions, come from coal and gas power plants [79]. Therefore, it is critically necessary to collect and sequester carbon for power plant emissions [80]. For efficient carbon dioxide removal, a coal power plant's exhaust gas should be discharged at ambient pressure [81]. Research on CO₂ capturing and keeping in storage is now being done to lower carbon dioxide pollution levels [82].

The three primary methods for capturing and separating carbon dioxide from CO₂/N₂ mixed gas are membranes, adsorption by porous materials, and liquid fluid absorption. Liquid materials, such as amine solutions [83], which were extensively used in industry, were frequently utilized for capturing carbon dioxide. Liquid sorbents have the benefit of being more successful at collecting carbon dioxide and under atmospheric circumstances, the watery basic fluid can preferentially capture and combine with acidic carbon dioxide gas [82]. However, the equipment's life span is impacted by the acidic nature of amine chemicals. Liquid sorbents can chemically interact with other corrosive gases in the gas mixture, which has the side effect of absorbing less active carbon dioxide [84]. In addition, the long-term cycling of liquid sorbents results in elevated regenerative temperatures as well as reactive response and heat deterioration [84, 85]. In other terms, this technique needs more energy because it typically involves regenerating aqueous solution at temperatures above 100 ºC [82].

Membranes have additionally been mentioned as an effective technique for separating carbon dioxide due to their high sensitivity; low energy needs, and ease [79]. Although membranes are effective at separating bulk carbon dioxide at high pressures, this technique is ineffective for separating carbon dioxide in low amounts because it requires more energy to squeeze the input gas [80]. As a result of the aforementioned disadvantages, many scholars are investigating more effective or appropriate solid adsorbents for carbon dioxide [86]. Solid adsorbents such as amine-supported silica [87], carbonaceous materials, zeolites [88], and metal–organic frameworks [89], have become more common than liquid absorbents as a result of their greater adsorption ability and cheaper price [84]. Adsorbents must perform the following characteristics satisfactorily: strong selectivity for carbon dioxide, high adsorption rates, and high regenerability [82].

Zeolites with consistent micropore sizes, excellent chemical stability, and large surface areas have been utilized in the exhaust gas separation industry among these materials [90]. However, the restriction of mass transfer resistance is a typical issue that many industrial zeolites encounter [91, 92]. Some methods have been put forth for reducing the barrier to gas molecules diffusing through zeolite pore channels: synthesizing zeolites with more substantial micropores [93], smaller compact particle sizes [94], and an open pore network within the solid's interior [95] are the first three methods. Mesopore formation depends on the use of expensive template agents, acid, or base post-treatments that may harm the crystal structure or result in pollution issues in the surrounding environment. Both methods (one and two) have a decent chance of finding commercial use because they are reasonably easy and safe for the ecosystem. It is well known that zeolite efficacy can be impacted by crystal size and shape because they change the typical passage time through the crystal [96,97,98,99].

Zeolite materials are commonly used for carbon dioxide removal, and CO₂ has a stronger preference for NaA zeolite than for NaX or NaY. Zeolite A's high selectivity is essential for removing carbon dioxide after burning because it allows for the separation of carbon dioxide from N₂ in the flue gas. Indira and Abhitha [100] investigated the CO₂, H_2, and N₂ adsorption capacities of Zeolite Na-4A. Zeolite A's carbon dioxide adsorption process involves both physical and chemical absorption. At low temps, physical binding is encouraged, and as the temperature increases, these materials function more difficult [100]. Large quadrupole moments of CO₂gas combine with the electric field close to the structure of the zeolite, leading to physical adsorption. High persistent quadrupole moment gas molecules can combine significantly with the zeolite-induced gradient of the electric field [101]. Carbon dioxide has a greater quadrupole moment than carbon monoxide, nitrogen, hydrogen, and methane (CH₄ < H₂ < N₂ < CO < CO₂) [102]. Carbon dioxide has a quadrupole moment of − 1.43 × 10¹³ cm², which is 3 times greater than N₂ [103]. Zeolite A therefore has a greater affinity for carbon dioxide than for N₂ or CH₄. Low Si/Al zeolites like X, Y, and A, also chemisorb carbon dioxide by producing bicarbonate or carbonate [104] in addition to the foregoing. The adsorption of carbon dioxide can be affected by pressure and temperature. The potential for carbon dioxide adsorption rises with an increase in carbon dioxide relative pressure and falls with a temperature rise. This is because at low pressures, the amount of carbon dioxide absorbed is exactly proportionate to the cationic density of zeolite, and at high pressures, the adsorption is governed by the number of zeolite pores [105].

Diffusion of carbon dioxide in zeolite 4A and H-ZSM-5 is forecast to rise with an increase in partial pressure [106, 107]. In addition to temperature and pressure, the zeolite framework structure's basicity and electric field have a significant impact on how well they function. Zeolites' basicity and electric field are caused by the type of counter-cation or exchangeable cation that is found in the framework holes of zeolites. As the basicity rises, zeolite A's ability to adsorb carbon dioxide grows. Li⁺ is the most basic cation, and it contributes to the basic character of zeolites [108], which have high carbon dioxide absorption capacities [109]. Si/Al ratio and particle size are some of the additional variables influencing carbon dioxide adsorption capability. The basicity and electric field of zeolites are inversely correlated with the Si/Al ratio. Zeolite A falls under the low Si/Al ratio group and has greater basicity as a consequence. This increases carbon dioxide adsorption. Another important consideration is the size of the pores, which need to allow adsorbate entry [110]. To guide the shape of the nanocrystalline NaA zeolite, Shakarova et al. [111] used methylcellulose. Due to the synthesis of zeolite nanoparticles with a 100 nm dimension, the product produced had improved carbon dioxide adsorption and a quick uptake rate [111]. Zeolite 5A is another excellent option for carbon dioxide adsorption, and studies have shown that it can reach equilibrium carbon dioxide adsorption rates of 3.38 mol/kg at 303 K and 100 kPa while only adsorbing 0.22 mol/kg of nitrogen under the same conditions [112].

The adsorption procedure of a zeolite structure was detailed in Fig. 4 by Roussanaly et al. [113]. On a zeolite surface with micropores, carbon dioxide produced from a source is specifically adsorbed [113]. The zeolite's ability to adsorb carbon dioxide depends on how it was synthesized [114]. Compounds made up of metal ions and groups with a specific structure that is considered are known as metal–organic frameworks [115]. They frequently have pores, and some organic compounds are frequently used as linkers or supports [116].

Fig. 4

Illustration depicting the structural composition of Zeolite A [114]

In the course of the elimination process, the voids within a material's structure often exhibit a solid nature, offering a unique opportunity for these gaps to be occupied by other desirable elements derived from the targeted substance, as elucidated by Willis et al. [115]. This intriguing phenomenon adds a layer of complexity to the elimination dynamics, introducing the possibility of incorporating beneficial constituents into the material's framework [115]. Extensive research has identified metal–organic frameworks (MOFs) as particularly well-suited materials for the capture and storage of various gases, including hydrogen (H₂) and carbon dioxide. The pivotal role of MOFs in gas sorption has been underscored by Abu Ghalia and Dahman [116], signifying their potential to address critical challenges related to gas storage and separation. The inherent adaptability of MOFs in accommodating and selectively interacting with different gases positions them as promising candidates for applications such as gas storage, separation, and capture. Their tunable structures, characterized by metal nodes and organic linkers, offer a versatile platform for tailoring the material to specific gas sorption requirements [116]. This versatility, coupled with the intriguing phenomenon of holes in the material being amenable to hosting desirable elements during the elimination process, further highlights the multifaceted nature of MOFs in addressing challenges related to gas management and environmental concerns.

Krachuamram et al. [117] conducted a comprehensive study on carbon dioxide adsorption using zeolites Na-X, denoted as Z₁ and Z₂, employing a combination of sodium silicate solution, aluminum hydroxide with sodium silicate granules, and sodium aluminate as silicon and aluminum sources. The characterization of the synthesized zeolites through XRD, BET, and surface studies unveiled that the utilization of sodium silicate and sodium aluminate sources resulted in the development of zeolite structures with notably large surface areas and cavity volumes. The carbon dioxide adsorption isotherms for Z₁ and Z₂ samples are illustrated in Fig. 5a and b, respectively. Intriguingly, the synthetic Z₁ samples exhibited a conspicuous reduction in carbon dioxide adsorption compared to the Z₂ samples. Further analysis of carbon dioxide adsorption sequences within the Z₁ samples revealed a hierarchy as follows: Z_{1CH_1D} < Z_{1CH_3D} < Z_{1_7D} < Z_{1CH_7D}. On the other hand, the carbon dioxide adsorption sequence for Z₂ samples was identified as Z_{1_7D} < Z_{1CH_7D} < Z_{1_3D} < Z_{1CH_1D} < Z_{1CH_3D}. Notably, Z_{2CH_3D} demonstrated the highest carbon dioxide adsorption potential at 4.25 mmol g⁻¹, coupled with the largest surface area and pore volume. Figure 6a and b presents plots depicting carbon dioxide adsorption with surface area and pore volume for Z₁ and Z₂ samples, respectively. The most remarkable finding was the exceptional performance of Z_{2CH_3D} as an absorbent, displaying a substantial capacity for adsorbing carbon dioxide (5.08 mmol g⁻¹ at 303 K) and exhibiting a relatively high affinity for carbon dioxide over nitrogen. This suggests the promising potential of Z_{2CH_3D} as an adsorbent in flue gas applications, indicating its capability to selectively remove carbon dioxide from gas mixtures [117].

Fig. 5

a Z₁ Isotherms for the absorption of carbon dioxide b Samples of Z₂ at 298 K [117]

Fig. 6

Examining the correlation between carbon dioxide adsorption, surface area, and pore capacity in samples a Z₁ and b Z₂ [117]

This observation highlights the capacity to augment surface areas and pore sizes as a means to concurrently elevate carbon dioxide adsorption efficiency. The study conducted by Krachuamram et al. [117] underscores that the effectiveness of carbon dioxide adsorption is intricately linked to the porosity volume and surface area of the adsorbent material. By systematically manipulating the synthesis parameters, the researchers were able to enhance both surface areas and pore sizes in the zeolite structures. This nuanced approach resulted in an evident improvement in the materials' ability to adsorb carbon dioxide. The direct correlation between the observed enhancements in carbon dioxide adsorption potentials and the variations in porosity volume and surface area emphasizes the critical role of these structural characteristics in dictating the adsorption performance of the zeolite materials [117]. This finding holds significant implications for the design and optimization of adsorbent materials tailored for carbon dioxide capture. It suggests that strategic adjustments in the synthesis process to achieve specific porosity volumes and surface areas can lead to materials with heightened carbon dioxide adsorption capabilities.

Boycheva et al. [118] conducted a comparative study on carbon dioxide adsorption using coal fly ash zeolites (CFAZ), specifically Na-Ca-X and Na-X. Zeolite X, a widely accessible product, is a frequently investigated absorbent for post-combustion carbon capture devices. This zeolite is characterized by a unique molecular structure comprising super-cages, resulting in a highly developed surface and large hole width surpassing the dimensions of the carbon dioxide molecule. In their investigations, Boycheva et al. [118] found that CFAZ exhibited a carbon dioxide absorption capacity comparable to their purified counterparts, despite having a smaller specific surface area. This intriguing phenomenon can be attributed to the greater number of adsorption sites in CFAZ, stemming from their larger unsaturated surface [118]. Additionally, the presence of iron species, distributed across the CFAZ surface and transferred from the ash, contributes to an additional increase in adsorption [119].

Dynamic carbon dioxide absorption patterns on the investigated coal fly ash zeolites, both in the absence and presence of water, are illustrated in Fig. 7a and b. Under ambient conditions, the molecular network of zeolites is influenced by polar water molecules that occupy the available holes with a potent attraction. This interaction between carbon dioxide and water molecules is a critical aspect of the CO₂/H₂O system, necessitating a nuanced understanding of the various adsorption types [118]. This insight is pivotal when evaluating zeolites as potential adsorbents for carbon dioxide capture applications.

Fig. 7

Deposition of carbon dioxide onto CFAZ breakthrough Curves: in a carbon dioxide flow (a); b Competitive adsorption in the CO₂/H₂O system [118]

The study not only sheds light on the comparable carbon dioxide absorption capacities of CFAZ and their purified counterparts but also emphasizes the intricate interplay between surface characteristics, additional adsorption mechanisms involving iron species, and the dynamic nature of adsorption patterns in the presence of water. These findings contribute valuable insights to the ongoing research on optimizing zeolite-based materials for efficient carbon dioxide capture and sequestration.

Within the structure of CFAZ_{M3_3}, the presence of more mesopores influences the mass transfer zone, resulting in a prolonged process, as depicted in Fig. 7a. A comparative analysis of adsorption kinetics reveals that CFAZ_AES, CFAZ_{M3_1,} and CFAZ_{M3_2} exhibit similar behaviors, while CFAZ_{M3_3} showcases an adsorption capability nearly twice as high as CFAZ_{M3_1} and CFAZ_{M3_2}, with CFAZ_AES ranking second in terms of adsorption efficiency [118]. This intriguing difference suggests that the carbon dioxide adsorption process under dynamic conditions is kinetically restricted and relies on the contact duration, emphasizing the role of interactions in improving adsorption.

In the context of pressure-swing adsorption devices, the higher adsorption capability observed at elevated pressures, especially in CFAZ_{M3_3}, proves advantageous for carbon dioxide capture. Despite CFAZ_{M3_3} having a slightly smaller surface area (421 m² g⁻¹) compared to CFAZ_AES (486 m² g⁻¹), it surpasses the latter in dynamic adsorption capacity. This discrepancy is attributed to the prevalence of Ca²⁺ ions in CFAZ_{M3_3}, demonstrating a greater capacity for carbon dioxide uptake than Na⁺ ions [120]. Addressing the challenges of using industrial zeolites for selective carbon dioxide removal in wet environments, it is recognized that zeolites tend to lose their adsorption characteristics in the presence of water due to their higher affinity for water molecules [121, 122]. To mitigate this, dehumidification is deemed necessary before carbon dioxide adsorption, albeit at the cost of increased energy usage, process complexity, and equipment expenses. Encapsulating zeolites within a casing emerges as a viable solution to limit water molecules' entry into the zeolite center, thereby preserving their carbon dioxide adsorption ability. Considering dynamic reaction-induced delay behavior for moisture, strategies involving amines for carbon dioxide removal have been developed. For instance, Liu et al. [123] introduced core–shell zeolite 5A@mesoporous silica-supported-amine (5A@MSA) hybrid adsorbents, enhancing carbon dioxide removal from exhaust gas in the presence of water [123]. The incorporation of a hydrophobic zeolitic imidazolate framework (ZIF)-8 layer further mitigates the impact of wetness on carbon dioxide adsorption processes, as demonstrated by Gao et al. [124]. This multifaceted approach involving encapsulation, amines, and hydrophobic layers showcases the innovative strategies employed to overcome challenges associated with water interference in zeolite-based carbon capture technologies [124].

Akisanmi [125] delved into the synthesis of zeolites derived from kaolinite and explored their efficacy in carbon dioxide removal methods. The study accentuates the potential of clay rocks, abundant and cost-effective raw materials, which exhibit remarkable versatility in the realms of adsorption and catalysis within various industries [125]. Utilizing metakaolinite as a precursor, the researchers successfully developed type A zeolite through chemical processes in alkaline solutions [126]. The choice of clay materials for zeolite synthesis is particularly intriguing due to their wide availability and economical nature [127].

Kaolin clay emerges as an ideal substrate for synthesizing type A zeolite, primarily attributed to its nearly equivalent molecular Si/Al ratio of one [128]. Kaolinite, a stratified aluminosilicate material with the molecular formula Al₄Si₄O₁₀(OH)₈ [129], belongs to the phyllosilicate group. Its use as a precursor in zeolite synthesis offers a sustainable alternative to purified chemical compounds like sodium silicate and sodium aluminate in conventional industrial processes [130]. This investigation emphasizes the significance of leveraging naturally occurring clay resources for zeolite production, highlighting both their economic advantages and environmental sustainability.

4.1.1 Carbon dioxide (CO₂) capture from N₂ and CH₄ mixture

Wu et al. [131] conducted a comprehensive investigation exploring the application of zeolite L, characterized by various crystal morphologies, for the adsorption of carbon dioxide from a mixture of CH₄ and N₂. The synthesis of zeolite L was approached with three distinctive crystal morphologies: circular (C-L), disk-shaped (D-L), and nanosized (N-L), each offering unique structural attributes. The study delved into the impact of potassium hydroxide (KOH) in the synthesis process, particularly in the formation of nanosized zeolite. In the synthesis of nanosized zeolite (N-L), the inclusion of KOH was found to be pivotal. A detailed examination revealed that when the molar ratio of K₂O reached 5.5, it led to the formation of N-L characterized by small crystals and a substantial crystallinity. This observation underscores the significant influence of KOH in tailoring the crystal size and morphology of zeolite L, specifically favoring the development of nanosized structures [131]. This research not only contributes to the understanding of zeolite L's applicability in carbon dioxide adsorption but also highlights the significance of crystal morphology and synthesis parameters in tailoring the material for optimal performance. The insights gained from this study could inform the design and optimization of zeolite-based materials for effective carbon capture from complex gas mixtures.

4.2 Measurements of FTIR adsorption

The characterization of the synthetic zeolite involved a comprehensive analysis using various techniques, including X-ray diffraction, Fourier Transform Infrared Spectroscopy (FTIR), N₂ adsorption, and Nuclear Magnetic Resonance of Solids. This multi-faceted approach aimed to provide a detailed description of the structural and functional properties of the synthesized zeolite. The carbon dioxide adsorption performance of the synthetic zeolite, conducted at 100 ºC for 48 h, demonstrated striking similarities to that of the commercial zeolite. This observation not only validates the efficacy of the synthesized zeolite but also positions it as a promising alternative to commercially available counterparts in carbon dioxide capture applications. To assess the durability of the zeolite, the researchers conducted multiple carbon dioxide adsorption cycles and examined the structural integrity using FTIR. The sustained presence of comparable bands across all cycles suggests that the zeolite retains its structure and carbon dioxide adsorption capacity. Remarkably, the FTIR analysis revealed a distinctive band at 2350 cm⁻¹ in the spectrum after 30 min of outgassing, with reduced intensity compared to Fig. 8 (the three cycles) [132]. This subtle difference may indicate some changes in the zeolite's surface or adsorption characteristics during the outgassing process, warranting further investigation.

Fig. 8

Bands for zeolite performed at 100 ºC and conducted at 48 h carbon dioxide adsorption at 20 ºC, outgassing, and three cycles of heating at 150 °C, 200 °C, and 300 ºC [132]

It implies that a smaller amount of carbon dioxide was physisorbed during this test, most likely as a result of small changes to the outgassing period. However, a similar pattern is seen because the carbon dioxide that has been chemically absorbed as monodentate, bidentate, or bridging carbonate species is highly adsorbed and needs the temperature to be raised to encourage its release and utilization in the following cycle [132].

4.3 Adsorption isotherm and selectivity in the ideal adsorption solution theory (IAST)

The substantially elevated adsorption of carbon dioxide compared to methane or nitrogen can be attributed to the higher polarizability and quadrupole moment of CO₂, as elucidated by Yu et al. [133]. This pronounced difference in adsorption performance is a consequence of the unique molecular properties of CO₂, which facilitate a stronger interaction with adsorbent surfaces. Figure 9a, b, and c illustrates a notable increase in carbon dioxide uptake across all samples, especially at lower pressures. This observation underscores the enhanced affinity of the studied adsorbents for carbon dioxide, emphasizing the potential for efficient carbon capture from gas mixtures [133].

Fig. 9

Adsorption and desorption of (a) carbon dioxide, (b) methane (c) Nitrogen and the selectivities for binary mixtures in adsorption of (d) N_2/CO₂(85:15 v/v), (e) CH₄/CO₂ (60:40 v/v), and (f) N₂/CH₄ (80:20 v/v) at 1 bar and 298 K for samples, N-L,D-L and C-L [131]

The intensified surface-carbon dioxide contact, crucial for effective adsorption, can be ascribed to the reduction in particulate size. As distinguished by Zarshenas et al. [134], the decrease in particle size contributes to an increase in surface energy, fostering a more robust interaction between the adsorbent surface and carbon dioxide molecules. This phenomenon is particularly significant at low pressures, where the surface energy becomes a determining factor in the adsorption process [134]. The molecular characteristics of CO₂, including its higher polarizability and quadrupole moment, confer a distinct advantage in terms of adsorption over methane and nitrogen. The observed increase in carbon dioxide uptake, especially at lower pressures, aligns with the enhanced surface-carbon dioxide contact facilitated by a reduction in particulate size.

It is compatible with the BET study that the carbon dioxide absorption capacity of sample N-L (70.7 cm³/g) was 20.9% greater than that of sample C-L (58.5 cm³/g). The anomalous rise in the absorption of carbon dioxide, methane, and nitrogen on sample D-L can be imparted to the addition of Na⁺, which increases the electric field force within the molecular sieve structure [135, 136]. To assess the possible separation abilities of the adsorbents, the IAST selectivities based on single-component isotherms were evaluated. Figure 9d, e, and f summarizes the IAST selectivity results for binary mixes of N₂/CO₂ (85:15 v/v), CH₄/CO₂ (60:40 v/v), and N₂/CH₄ (80:20 v/v) on samples C-L, D-L, and N-L. Similar IAST selectivities for CH₄/CO₂ (60:40 v/v) were presented by N-L,D-L, and C-L, with values of 75.3, 75.7 and 72.9, respectively. The IAST selectivity N-L (198.6) was considerably greater than those for C-L (188.6) and D-L (181.8) for the N₂/CO₂(85:15 v/v), demonstrating that the prior can successfully separate carbon dioxide and nitrogen [131].

Commonly used in commercial N₂/O₂ separation operations, lithium-aided low silica X type (Li-LSX) zeolite is well established, but its effectiveness for carbon dioxide was unknown at the time. Using a packed bed setup, the power of Li-LSX was evaluated for carbon dioxide capture and CO₂/N₂ selectivity in the post-combustion state in the presence of 14% carbon dioxide. It was determined to be 4.43 mmol/g and the ratio to be 85.7 at 60 ºC [137, 138]. The selectivity originally decreased slightly as the temperature increased but at calcination temperature, it increased to 128.1 and the adsorption rate even increased due to the effective rise in basicity. The Avrami kinetic model was used to determine the co-existence of different adsorption processes. Additionally, it is implied that Le Chatelier's principle will cause a rise in partial pressure to accelerate adsorption. The reducing of post-combustion carbon is therefore possible using this potential sorbent. Li-embedded zeolites eventually even demonstrated their capability for carbon adsorption and storing in natural environments [139]. Zhao et al. [136]provided dynamic data in 2018 that showed the influence of Li⁺ over Na⁺ as it interacted with ZSM-25's highly selective adsorption capability. A change in alkali size increased the rate of adsorption by 6.1% (303 K & 9.5 pressure), which resolved the issue of commercial application in selective CO₂/CH₄ capture in natural gas cleaning and biogas upgrading [136].

The LSX zeolite was subsequently changed by the Yang group to contain commercially feasible strontium (Sr), with encouraging findings for purified Sr-LSX zeolite in air separation [140]. The enhanced findings serve as proof that, among all adsorbents, the newly formed LiX-80 zeolite has a hopeful CO₂/N₂ selectivity. By adding palladium (II) and silver (I) to LiX-80, a new polymetallic cation swapped zeolite-LiPdAgX was developed based on this finding. In comparison to 13X and LiX-80 zeolites, it displayed a greater level of carbon dioxide and increased CO₂/N₂ selectivity. Confined fluidization has been shown to increase the efficiency of the adsorption process, making it a viable option to the conventional approach for CO₂ adsorption using a fluidized bed of 13X zeolite particles [141]. Due to carbon dioxide's molecular durability and non-reactive character, there are still difficulties with removing carbon dioxide, especially in the lower pressure area below 1000 ppm. Amazingly, it was discovered that Ba(II) swapping MFI-type zeolites provided a successful practical and theoretical solution to the difficult activity. It was revealed by IR and X-ray diffraction that a bridged M²–O²⁻–M²⁺ shape is the base of the stable carbon dioxide binding (C–O interaction as well as O–Ba²⁺ interaction) [142], and this explanation illustrates in great detail why the extraordinary characteristic still holds at 300 K. Using Na atoms, Pulido et al. [143]thoroughly demonstrated the safety of the bridge using the computation [143].

Researchers, Cheng et al. [144] have examined the testing adsorption of carbon dioxide onto amines of various types applied to HZSM-5 zeolites. This adsorbent was synthesized by wet impregnation technique to attach amines to HZSM-5 zeolites. After investigating the carbon dioxide adsorption characteristics in fixed-bed adsorption technology, structural, thermal, and kinetic studies were conducted. In comparison, carbon dioxide adsorbents supporting monoethanolamine (MEA) and hydroxyethyl ethylenediamine (AEEA) demonstrated good carbon dioxide adsorption capacities, with maximum values of 4.27 and 4.44 mmol/g being achieved, respectively. The HZSM-5, with a 25 Si/Al ratio and 2.3 m in average particle size, was found to be a favorable carbon dioxide adsorbent support. The carbon dioxide desorption process of AEEA-embedded HZSM-5 indicated a low activation energy of 54.27 kJ/mol, according to desorption rates and thermodynamic investigations [144]. Table 2 summarizes the capability of various kinds of zeolites to absorb carbon dioxide.

Table 2 Summary of carbon dioxide capture capacity of different types of zeolites

This low activation energy may be attributed to two factors: first, the HZSM-5 zeolite offers a significant amount of free H⁺ ions, which directly contribute to carbamate breakdown; second, the HZSM-5 zeolite might also contain Al atoms, which can attach to the carbamate's N atom and stretch the C-N bond, thereby facilitating decomposition. The regenerative cycle experiments also showed that the AEEA-embedded HASM-5 adsorbent has high adsorption ability, demonstrating the energy effectiveness of the carbon dioxide desorption process and cheap carbon dioxide capture expense [144]. Although microporous zeolite A provides effective CO₂ adsorption, the tiny pore size results in diffusional and mass transfer resistance. To attain improved CO₂ adsorption characteristics, hierarchical zeolite A, a zeolite with customized molecular and physical traits, must be developed.

5 Sulfur dioxide (SO₂) removal using zeolites

One of the most crucial environmental issues in the developing world is atmospheric pollution [165]. Due to the resulting acid rain and lung damage that is extremely dangerous to environments and people, sulfur dioxide, which is primarily generated through the use of natural fuels containing sulfur has recently drawn a lot of focus [166]. For the elimination of sulfur dioxide, wet scrubbing, and dry sorption techniques are part of the flue gas desulfurization technology. However, the primary drawbacks of these techniques are their high prices, secondary contamination, and poor effectiveness in the wet process and dry process, respectively [167]. Researchers have recently paid a lot of attention to the use of photocatalytic processes for the elimination and scrubbing of atmospheric pollutants. As a result of its non-toxicity, molecular stability, simplicity of production, and environmental kindness, TiO₂ is regarded as the most encouraged photocatalyst [168]. However, TiO₂ has many disadvantages that restrict its use in photocatalytic processes, including the ability to only capture UV light [169], the fast recycling of electron–hole pairs, and limited surface area [170].

The growth of solar energy as a sustainable and pure energy has more favorable effects on the economy and the natural environment, even though the light source for photocatalytic processes could either be created or solar energy [171]. However, only 5% of solar energy is in the UV spectrum, while 45% of solar energy is in the visible spectral range. As a result, purified TiO₂ has a restricted ability to photocatalyze the degradation of atmospheric pollutants when exposed to sun light [172]. It has been stated that the constraints of TiO₂ can be overcome by doping, semiconductor coupling, and the use of compounds based on TiO₂ [173]. The integration of adsorption and photocatalytic processes has led to the recognition of the composites as appropriate materials for the elimination of pollutants. The harmful substance may be absorbed by the absorbent and then distributed to the photoactive spots in a mixture [174]. Additionally, the movement of undesirable by-products into the adsorbent may replenish the active sites and increase the photocatalyst's lifespan. Recently, there has been increased interest in the use of photocatalyst compounds in gas removal devices [175]. Due to the practical challenges of holding big gas quantities, photo degradation of air pollutants is typically only performed at a single photo reactor [176]. Utilizing effective photocatalysts with high absorption of gas efficacy could be a suitable option to get passed this technological barrier [177].

5.1 Sulfur dioxide removal using Tio₂-Ze composite

5.1.1 Effect of gas temperature and sunlight intensity

Figure 10a and b show the outcomes of two common testing methods to determine the adsorption including the photocatalytic performance of TiO₂-Ze composite for the elimination of sulfur dioxide in atmospheric conditions during 10 h in the summer from 7.30 a.m. to 17.30 p.m. These investigations were performed to identify adsorption or desorption mechanisms associated with the temperature rise caused by radiation from the sun. A slight increase in the output content of sulfur dioxide was observed in the first investigation on the adsorption efficacy in atmospheric circumstances during the night. As shown in Fig. 10a, the elimination of sulfur dioxide decreased statistically significantly from 22.7 to 16–17% (p < 0.05) as the atmospheric temperature rose from 22 to 48 °C and the gas temperature rose from 17 to 44 ºC. The adsorption of sulfur dioxide on the composite was then comparatively enhanced with the drop in atmospheric temperature (p < 0.05). Since the air or gas temperature hindered the adsorption activity, the morning and afternoon were the optimum times to investigate the adsorption performance of sulfur dioxide in atmospheric circumstances [171, 178].

Fig. 10

TiO₂-Ze composite removal of sulfur dioxide daily time (a) adsorption in dark conditions, (b) Adsorption operations (50.0% TiO₂-Ze at 300 ºC, SO₂ initial concentration: 8 ppm, composite dosage of 0.1 g, and gas flow rate of 40.8 mL min⁻¹) [179]

To quantify SO₂ photodegradation on a combination of TiO₂-Ze in atmospheric conditions, another experiment was conducted. When the intensity of the light intensified from 87 to 110.0–140.0 mWcm⁻² over a day, the photocatalytic process' ability to remove sulfur dioxide improved, and the highest photodegradation of sulfur dioxide (34–35%) was noted between the hours of 10:30 and 14:30 (Fig. 10b). The efficacy of the photocatalytic process for the elimination of sulfur dioxide gas was substantially reduced to 25–28% before and after these periods (p < 0.05). As a result, there are significant variations in the elimination of sulfur dioxide during times of rising or declining sun radiation. The temperature in the photocatalytic reaction was overcome by the influence of the increased sun energy durability. In reality, as incoming light strength rises, more photocharges are produced by the photocatalyst, which speeds up the decomposition of pollutants [180]. Based on the findings, the following evaluates for sulfur dioxide elimination were carried out at temperatures in the range of 40–44 ºC and the highest sunshine strength (125.0–135.0 mWcm⁻²). The adsorption of sulfur dioxide on the polymer could take place chemically or mechanically [181]. The van der Waals force between sulfur dioxide molecules and the two substances could be the cause of the physical binding [182]; oxygen holes on the surface of the TiO₂ [183]; and consistent microporous canals on the zeolite surface [165].

TiO₂ and zeolite have polar groups on their surfaces that can adsorb the polar SO₂ molecule and are primarily responsible for the process of chemical adsorption [182, 184]. Zeolites and TiO₂ have surface OH groups that may form H-bonds with sulfur dioxide, which is the main mechanism by which the adsorbate and adsorbents engage with each other [50, 185]. It is possible to hypothesize the following processes for the photocatalytic removal of SO₂ on TiO₂ when it is exposed to radiation as shown in reactions 1–7 [186].

$${\text{Ti}}{\text{O}}_{2 }+{\text{hv}}\to {\text{h}}_{\text{vB }}+{\text{ e}}_{{\text{CB}}^{-}}$$

(6)

$${\text{O}}_{2 }+{\text{ e}}_{{\text{CB}}^{-}} \, \to \, {{\text{O}}_{{2}^{-}}}$$

(7)

$${{\text{O}}_{{2}^{-}}}+{{\text{H}}^{+} \to \text{HO}}_{2}$$

(8)

Following that, holes and electrons confined the immobilized sulfur dioxide and oxygen into the active groups at the photocatalyst surface traps (see in reactions 6 and 7).The active oxygen that was seen in steps 8 and 9 could then oxidize SO₂⁺ to produce SO₃ [187, 188].

$${\text{SO}}_{{{\text{2(ads) }}}} + {\text{h}}_{{{\text{vB}}^{ + } }} \to {\text{SO}}_{{2({\text{ads}})^{ + } }}$$

(9)

$${\text{O}}_{{{\text{2(ads) }}}} + {\text{e}}^{\_} \to {\mkern 1mu} {\text{O}}_{{2({\text{ads}})}}$$

(10)

$${\text{SO}}_{{2({\text{ads}})^{ + } + }} {\text{O}}_{{2({\text{ads}})}} \to {\text{SO}}_{{3({\text{ads}})}} + {\text{O}}_{{({\text{ads}})}}$$

(11)

$${\text{SO}}_{{{\text{2(ads) }}}} + {\text{O}}_{{2({\text{ads}})}} \to {\text{SO}}_{{3({\text{ads}})}}$$

(12)

The adsorption capacity of various zeolites onto SO₂ has been summarized in Table 3.

Table 3 Summary of adsorption of SO₂ using different zeolites

6 Hydrogen sulfide (H₂S) adsorption from natural gas

The past century has seen several persistent economic and environmental issues related to the capture and reduction of hydrogen sulfide. Syngas, natural gas, biogas, landfill gas, coal furnace gas, refining gas, etc. all contain H₂S, a highly unpleasant, pungent, and poisonous substance that is also present in effluent streams. It is essential to remove this substance from these waterways for reasons of both economy and protection. Equipment and pipes corrode as a result of H₂S' innate propensity to generate a corrosive solution when exposed to water. Additionally, it poisons catalysts and lowers the burning value of the combustion vapors [200].

In natural gas, the concentration of H₂S ranges from 4 to 1000 parts per million (ppm), and it varies greatly depending on the location of the producing region [201,202,203]. Due to its danger to human health, ability to contaminate downstream catalysts, and ability to cause machinery to corrode, hydrogen sulfide is typically regarded as a major issue for the chemical industry [204,205,206,207]. So, to satisfy the needs of future operations, natural gas H₂S amounts should be kept below 2 ppm. By using the recently developed wetness method, Montes et al. [208] prepared the Cu/MSU-1 and Zn/MSU-1 adsorbents to remove H₂S. They observed that adding small amounts of Cu and Zn enhanced the removal ability to adsorb H₂S of MSU-1, and this showed the adsorption abilities of 19.2 mg g⁻¹(20Cu/MSU-1) and 42.3 mg g⁻¹(10Zn/MSU-1), respectively [208].

The effects of 13X on H₂S (50–100 ppm) removal at normal biogas operational conditions were investigated by Bareschino et al. [209]. Due to the production of sulfur and polysulfides, they deduced from the study of the concentration patterns that the system never achieves full concentration. They also observed that a trace quantity of water facilitated the adsorption of H₂S [209]. The adsorption process originally used physisorption, but after saturation, it switched to chemisorption, according to temperature profiles from adiabatic experiments. Using a Cu-exchanged 13X zeolite sorbent (13X-Ex-Cu), Barelli et al. [210] produced a desulfurized fuel appropriate for liquid carbonate fuel cell (H₂S requirement of 1 parts per million) by capturing hydrogen sulfide from biogas (200–1000 ppm H₂S) [210]. The superior desulfurization efficiency of 13X Ex-Cu as compared to 13X investigated by Bareschino et al. [209] in comparable situations is due to the abundance of Cu²⁺ ions, which results in an effective physical–chemical adsorption. 13X Ex-Cu capabilities increased as temperature increased. In addition, a decline in breakthrough time and capacity was expectedly brought on by a rise in input sulfur concentration [209].

To assess the effectiveness of the resulting AgNaA nano-zeolite in eliminating H₂S from biogas for application in solid oxide fuel cells, The NaA nano-zeolite has been synthesized and enhanced by silver ions using the ion-exchange technique by Bahraminia et al. [211]. In contrast to conventional 4A and not modified NaA nano-zeolite, AgNaA nano zeolite showed a prolonged breakthrough time of 310 min and a greater capacity of 33.24 mg/g to reach 1 ppmv of H₂S in the output gas. These improvements were made by decreasing the crystallite size of the adsorbent and adding Ag⁺ ions to the zeolite structure [211]. The greater amount of water that developed in the AgNaA sample may be an indication that this adsorbent will work better due to a higher likelihood of H₂S chemisorption. AgNaA renewal results in a minor shift in adsorption capacity (about 5% per cycle), which is ascribed to zeolitic water loss brought on by elevated regeneration temperatures.

Grand canonical ensemble Monte Carlo (GCMC) models were run by Yan et al. [212] to assess the capacity of 95 different types of all-silica zeolites to remove the six poisonous gases SO₂, NH₃, H₂S, NO₂, NO, and CO. The modeling results demonstrated that zeolite structures with open surface areas of 1600–1800 m²g⁻¹ and pore diameters of 0.6–0.7 nm, such as PAU and AFY, effectively capture H₂S, NO, NO₂, CO, and NH₃. However, their findings demonstrate that the capability for NH₃ over H₂S is either the same or higher for both adsorbents. The best adsorbents for H₂S adsorption are found to be MER, PAU, and AFY with loadings of 5.4, 5.5, and, 7.8 mmol g⁻¹, respectively. Furthermore, the authors concluded that zeolites with a surface area of roughly 1700 m² g^-1 and a cavity percentage of roughly 0.3 are suitable H₂S removal prospects. However, it has been discovered that the zeolites ITT, TSC, JSR, IRR, and RWY have superior abilities to adsorb all of the gases. RWY has the greatest H₂S storing capability of 17.74 mmol g⁻¹, while the other 9 zeolites can adsorb H₂S between 8 and 10 mmol g⁻¹ [212]. The most suitable materials for H₂S adsorption have an available surface area of 800–1600 m² g⁻¹ and a pore width of 0.7–0.9 nm, which are roughly in the range of the values given by Song et al. [213] for different all silica zeolites. They noticed that at low pressures, the Van der Waals force predominates, and at high pressures, the space or hollow area of zeolites plays a crucial part. H₂S, CO, CO₂, H₂O, and N₂ were all present in the blast furnace gas input, and CHA had the greatest adsorption capacity with the best H₂S selectivity, while FAU and LTA had exceptional H₂S selectivities at low adsorption capacities [213].

Georgiadis et al. [214] used an industrial molecular sieve with a Si/Al ratio of 0.97, Ca and Na non-framework cations, and a structure approximating an LTA-type zeolite like 3A or 4A. Analysis based on temperature changes, activation energies, and thermodynamic studies point to exothermic and spontaneous physisorption. 15 effective regeneration rounds were completed with a capacity reduction that was within the margin of trial error. H₂S was able to reach its maximal equilibrium capacity of 193.9 mg/g at 1 atm and 25 ºC. However, as the supply gas's CO₂ concentration rises, so does its ability to adsorb H₂S. With a feed matrix containing 36% CO₂ and a feed gas containing 3000 ppmv H₂S, the equilibrium capacity dropped from 164.5 mg/g in a carbon dioxide-free feed matrix to 57.7 mg/g. The research's desorption phase was performed at 200 ºC in contrast to many other studies that carried out regeneration at temperatures higher than 350 ºC [214].

Another researcher, Zhang et al. [215], looked into a new adsorbent for removing hydrogen sulfide from natural gas that was assisted by nanoparticles on SBA-15 zeolite. By obtaining nano-sized active metal oxides through the mesopores of zeolites, this study offered a practical method for removing hydrogen sulfide from natural gas [215]. Equation 13 can be used to determine the efficiency of the synthesized material for hydrogen sulfide removal [216].

$$\text{Q}=\frac{({C}_{in}-{C}_{out})}{{C}_{in}}*100\%$$

(13)

η: the effectiveness of removing hydrogen sulfide; C_in: the hydrogen sulfide concentrations at the intake (ppm); C_out: the hydrogen sulfide concentrations at the exit (ppm).

To determine the hydrogen sulfide removal capability of the adsorbents, stability evaluation processes for H₂S removal also took place at 30 °C and 30,000 h⁻¹. When H₂S removal rates attained 90%, the capture capacity was calculated as the mass of adsorbed hydrogen sulfide divided by the mass of adsorbent per unit time. The equation is as follows [217].

$$\text{Q}=\frac{{LM}_{s}*({C}_{in}t-{\int }_{0}^{t}{C}_{out}dt)}{1000m{V}_{L}}$$

(14)

Q: The hydrogen sulfide adsorption capability (mg g^-1); C_in: The H₂S input amounts (ppm); C_out: the fixed bed reactor's output H₂S amounts (ppm); Ms: the hydrogen sulfide molar weight (34 g mol⁻¹); t: the duration of the adsorption (min); V_L: the molar volume of the gas (24.45 L mol⁻¹); m is the adsorbent mass (g), and L is the gas discharge rate (L min⁻¹).

At 30,000 h⁻¹ and 30 degrees Celsius, the synthesized 1Cu1Zn/SBA-15 adsorbent demonstrated remarkable H₂S removal effectiveness, and the H₂S capture capacity attained 296.78 mg g⁻¹ at a 10% breakthrough [217]. Table 4 illustrates the adsorption of H₂S with various zeolites.

Table 4 Summary of H₂S adsorption using Zeolites

7 Economic analysis

Presently, the global expense for capturing CO₂ is between 60–110 USD/t, with projections indicating a decrease to 30–50 USD/t by 2030. This cost reduction will enhance the adoption of these technologies on a large scale in industrial settings [225].

The economic evaluation of a four-step vacuum swing adsorption (VSA) method for post-combustion CO₂ capture was carried out by Subraveti et al. [226]. There was 20 mol% CO₂ in the steam-methane reformer's dried flue gas. The VSA method was optimized to minimize CO₂ capture costs using Zeolite 13X, which has requirements of 95% CO₂ purity and 90% CO₂ recovery. According to the findings, each metric ton of CO₂ eliminated using the VSA method cost 90.9€ in expenditures. The capital costs for the VSA process came to 37% (18.6€/metric ton CO₂) of the avoided cost of CO₂, with contributions from vacuum pumps (8%/metric ton CO₂ avoided), compressors (about 3%/metric ton CO₂ avoided), and columns (8%/metric ton CO₂). Conversely, operating expenses, which comprised both fixed and variable operating costs, made up roughly 62%. The electricity consumption of 18.15€/metric ton CO₂ averted cost, or around 38% of the capture cost, made up the majority of the contribution. About 5.5% (1.95€/metric ton CO₂) of the overall capture expenses were made up of adsorber costs. Figure 11 shows the cost performances for the deployment of post-combustion carbon capture and storage in steam-methane reformer facilities. The CO₂ avoided costs component was displayed in euros per metric ton of CO₂ avoided. It is observed that the primary factor in the adoption of carbon capture and storage in both methods continues to be CO₂ capture cost. Figure 12 illustrates how adsorbent prices affect the lowest cost of CO₂ capture. Each data point on the graph corresponds to a distinct optimization run, providing a comprehensive view of the relationship between MOF price and the associated minimum cost of capture [226].

Fig. 11

Cost breakdown of the VSA technology using zeolite 13x, taking into account expenses for drying and cooling [226]

Fig. 12

Prices of adsorbents have an impact on Zeolite 13X's minimum CO₂ collection cost, which includes VSA drying and chilling expenses. Indicators 1 × , 2 × ,…, 10 × denote the multiplier that has been applied to the cost [226]

For coal-burning power plants, solar-assisted pressure–temperature swing adsorption (PTSA) CO₂ adsorption capture by zeolite 13X was recommended. The power plant produced 800MWe annually by operating 334 days a year. Either solar thermal energy or steam coming from the turbine provided the heat needed for CO₂ regeneration. For the carbon emission intensity (CEI) of 94 g/kWh, the economic analysis was conducted in terms of the levelized cost of energy (LCOE) and cost of CO₂ avoidance (COA). According to the study, the LCOE and COA of solar-assisted power plants were roughly $73/MWh and $39/tonCO₂, respectively [227].

In light of COA, an economic analysis of CO₂ adsorption using zeolite 13X emitted from a 500 MW power station was conducted. 13% of the content was CO₂ on a dry basis. According to the analysis, COA amounted to $51/ton of CO₂ averted [228]. PVSA was used to absorb CO₂ emissions from a 500 MW power station along with other flue gas emissions. A 90% purity and 90% recovery CO₂ adsorption was required. According to the economic evaluation, there was a $33.4 total annualized cost per tonne of CO₂ avoided [229].An economic analysis of a 550 MW power plant using zeolite 13X to adsorb CO₂ from flue gas with a 15% CO₂ content was conducted. It was discovered that with 95% CO₂ purity and 90% recovery, the COA was almost $26.3 per ton of CO averted [230].

The study examined the financial implications of employing Zeolite 13X for CO₂ post-combustion capture. The inlet temperature and pressure were 30 °C and 1.3 bar, respectively, while the flue gas stream included 12% CO₂ v/v and 95% relative humidity. 96% purity of CO₂ was required, and the factory processed 5000 tons of flue gas each day. At the flue gas stream outflow, the temperature and pressure were 30 °C and 110 bar, respectively. According to the study, the minimum CO₂ capture cost at 90% recovery for the TSAD configuration (six-step temperature swing adsorption) was almost 78 € per ton of CO₂captured. Figure 13 illustrates how different energy sources affect operating costs. As previously noted, the AIC consistently accounts for the greatest portion of the overall expenses (Fig. 13a). The expenses associated with the adsorption-based process primarily stem from the investment in adsorption columns (Fig. 13b), which escalates consistently with both the contactor's size and the quantity of units needed. For VSA and TSA, the cost share of the adsorbent was roughly 5% and 3%, respectively. Figure 13c shows how different energy sources affect operational costs. Electricity's higher price compared to steam increases the total AOC for the VSA9 configuration compared to TSA cases, while exergy consumption remains similar across all scenarios. While the energy consumption figures are similar in both circumstances, the higher cost of electricity relative to steam has an impact on the overall AOC of the VSA9 configuration as compared to those of the TSA configuration. Flue gas drying and CO₂ conditioning expenses make up 25% and 35% of the AIC and AOC, respectively [231].

Fig. 13

Costs per unit of CO₂ collection broken down for low-cost systems with 90% recovery; a annualized investment cost (AIC); b expenses associated with the adsorption-based process primarily stem from the investment in adsorption columns; c different energy sources that affect operational costs, TECH stands for technology for adsorption columns [231]

Using Zeolite 13X, the pressure-vacuum swing adsorption (PVSA) cycles' economic performance for post-combustion CO₂ capture was examined. To treat flue gas, 95% purity and 90% CO₂ recovery were needed. The minimal CO₂ averted costs rise as CO₂ compositions drop, as seen in Fig. 14a. For instance, the minimum avoidance cost increased from 12.2 to 138.5€ per ton of CO₂ as the content of CO₂ reduced from 30 to 3.5%.

Fig. 14

a the four-step PVSA cycle's minimum CO₂ avoided costs at various CO₂ compositions. b A comparison of the four-step and six-step dual reflex (DR) PVSA cycles with SO₂ Avoided costs (c) Breakdown of investment cost limits of the four step cycle (d) Breakdown of the six-step cycle's investment costs (CAPEX) related to cost limits of the six step DR cycle. e Operational expenses (OPEX) breakdown for the four-step cycle. f Operational expenses (OPEX) breakdown for the six-step DR cycle [232]

As illustrated in Fig. 14b, the minimum cost of CO₂ avoidance per metric ton of CO₂ is higher for the four-cycle PVSA compared to the six-cycle PVSA across the entire spectrum of CO₂ compositions. The capital expenditure for CO₂ avoidance per metric ton of CO₂ surged from €3.8 to €28.3 as the CO₂ composition in flue gases decreased from 30 to 3.5% in the four-cycle PVSA adsorption, as depicted in Fig. 14d. Electric power accounted for the majority of the operating costs across all CO₂ compositions in flue gases for the four-cycle PVSA, as shown in Fig. 14f. The capital cost shared between 24 and 32% of the CO₂ avoided cost, as illustrated in Fig. 14c and e for the range of CO₂ composition. The operating expense's share of the total cost increased from 68 to 76% when the CO₂ composition decreased from 30 to 3.5%. Operating costs ranged from 77 to 84%, with electricity usage being the main cause [232].

We investigate a pressure/vacuum swing adsorption (PVSA) method for CO₂ post-combustion collection from dry flue gas using zeolite 13X. Using a 4-step cycle with light product pressurization, we evaluated the performance of a PVSA-based process for a 500 MW power plant using a dry N₂-CO₂ gas mixture. The overall projected yearly cost of CO₂ capture from a dry N₂-CO₂ gas was $31.8 per metric ton of CO₂ avoided ($29 per metric ton of CO₂ caught), and $33.4 per metric ton of CO₂ avoided (US$30.4 per metric ton of CO₂ captured) from a wet N₂-CO₂ gas mixture [229].

8 Conclusion and challenges, and future outlook

This review paper investigated the CO₂, SO₂, and H₂S removal from biomass-derived flue gas. The elevated carbon concentration of the atmosphere has been regulated over time by using a variety of methods. Chemical absorption, membrane separation, and chemical cycling are some of the various methods for capturing the aforementioned gases. Zeolites have several advantages, including low corrosion, tolerable cavities, and pore sizes, high surface area, excellent thermal stability, low thermal expansion, ease of synthesis, high adsorption capacity, and selectivity, making them deserving of further study for potential use in gas removing, particularly CO₂, SO₂, and H₂S in a post-combustion process.

The Si/Al ratio, particle size, size of molecules adsorbed, and external factors like pressure and temperature are the main variables that affect zeolites' ability to adsorb carbon dioxide. Temperature has a significant correlation with pore size in zeolites and has a considerable effect on their capacity to adsorb CO₂. However, when the temperature rises over 100 °C, zeolites' ability to adsorb CO₂ drastically diminishes. In general terms, the optimal temperature for CO₂ adsorption using zeolites as adsorbents is 70 °C, which is slightly less than the actual flue gas temperature, which is usually 90 °C. As a result, the flue gas must be cooled or the adsorbents have to be modified to make them appropriate for high-temperature adsorption before such adsorbents may be utilized for the capture of CO₂ in industrial settings. According to the accounts to be presented, zeolite A has an adsorbent capability of about 6.12 mmol/g of carbon dioxide, and this can be improved with appropriate chemical modifications. Pore size modification and amine functional modification are two changes made to the zeolite framework structure that improve the material's capacity to adsorb carbon dioxide. Chemically modified zeolites with amine, silica, and ion exchange processes improve the adsorption of CO₂behavior. In the years to come, dual chemical alteration of zeolites will need to attract more attention since it provides an intriguing possibility for improving carbon dioxide removal performance. In this review, synthetic/natural zeolites that can be used for removal of SO₂ were also investigated. Under optimal circumstances, the zeolites had the highest SO₂ elimination efficiency of 78.8%. Overall, zeolites, whether natural or synthetic, could be used to remove SO₂ effectively. For zeolites to be an effective adsorbent for hydrogen sulfide removal, they must possess stable structures, high sulfur loading capacities, and good regenerability. According to this review, Cu1Zn/SBA-15 adsorbent, among synthetic zeolites, had the greatest H₂S elimination efficiency; the greatest H₂S capture capacity, a broad temperature of operation ranges, as well as excellent regeneration properties.

The zeolites that have been studied here have demonstrated efficacy in the adsorption of gaseous CO₂, SO₂, and H₂S; yet, to fully benefit from adsorption-based flue gas clean-up technology, these zeolites must first overcome many challenges. Zeolite's limited usable operating conditions, high synthesis costs, inability to remove numerous impurity gases at once, and lack of long-term stability are the main obstacles to its usage in flue gas clean-up. At very low pressures, zeolite isotherms often show a sudden increase in adsorption capacity, which is followed by a plateau. Therefore, if the desorption pressure is atmospheric pressure, it is challenging to establish good functioning capacities. To get good operating capacities in this situation, vacuum swing adsorption using zeolites is more appropriate. On the other hand, flue gas capture requires more energy when deep vacuum is used. Finding zeolites and zeotypes that exhibit a less steep adsorption isotherm at extremely low pressure would therefore be desirable. Zeolites that have undergone ion exchange, flexible zeolites, or zeotypes might all be used to do this. Work on the adsorption of flue gas from wet gas streams has increased in the past several years. This is difficult for zeolites with low Si/Al ratios because of their high affinity for water, which competes with flue gases for the adsorption sites. Higher Si/Al ratio zeolites are less susceptible to water and, therefore, more suited for operations involving wet gas streams; however, this comes at the expense of their adsorption capability. A thorough investigation into the effect of trace impurities on the lifespan and functioning capability of zeolites is still lacking. If zeolites are used extensively, this should be taken into consideration.

Stable zeolites still need to be developed, despite reports of reasonably high CO₂, SO₂, and H₂S adsorption capabilities of diverse zeolites. Zeolites may contain a variety of functional groups, such as MOFs, and combinations of hybrid materials with various functions may be developed for cooperative interactions. The prospect of simultaneous flue gas clean-up presents new problems for materials design since the use of multifunctional materials can potentially enable the removal of flue gases while also developing conceptually new integrated unit processes. Indeed, the introduction of new materials that can remove several flue gas pollutants (CO₂, SO₂, and H₂S concurrently) could result in a significant reduction in capital, operating, and energy needs. In the future, there will be an opportunity to employ these innovative materials in flue gas clean-up applications due to this knowledge gap.

To increase zeolites' selectivity, capacity, and adsorption rate at lower temperatures, further work needs to be done. The stability of zeolites over the long term is another crucial factor. Maintaining the zeolites' capacity during extended cycles of operation is crucial from an economic standpoint. Adsorbent materials must also be able to capture acid gases in the whole range of humidity conditions with limited degradation since flue gas contains considerable amounts of water. Furthermore, the adsorbent must be resistant to significant capacity loss in the presence of oxygen to function as an acceptable unit and effectively remove these acid gases. Consequently, further efforts are required to enhance the stability of zeolites in humid environments. Overall, there are an enormous number of options for developing and altering zeolites that effectively capture acid gases from flue gas streams.

Notably, we believe that a significant challenge lies in the search for zeolite with all the desired properties, such as a high capacity for H₂S adsorption, notable selectivity, and full regeneration capabilities, especially because materials that are not commercially available can often yield promising results.