INTRODUCTION
O VERVIEW IN THE FAU TYPE ZEOLITES , X ZEOLITES
Zeolites are crystalline aluminosilicates of alkali or alkali earth elements, such as sodium, potassium, and calcium, and are represented by the following chemical formula [1]:
Zeolites are represented by the formula M x/n [(A1O 2 ) x (SiO 2 ) y ] ã zH 2 O, where x and y are integers with y/x equal to or greater than 1, n denotes the cation M's valence, and z indicates the number of water molecules per unit cell The fundamental building blocks of zeolites are silicon and aluminum tetrahedra (SiO4 and A1O4), which combine to form secondary polyhedral structures such as cubes, hexagonal prisms, octahedra, and truncated octahedra These tetrahedra are interconnected by shared oxygen atoms, creating a three-dimensional crystalline framework This versatile arrangement allows for approximately 800 potential crystalline structures, although only about 200 have been identified in natural deposits or synthesized.
The framework for faujasite (FAU)-type zeolite (Figure 1.1A) can
7 be constructed by linking sodalite cages through double six-rings [6, 7, 32–
The FAU-type zeolites feature a significant structure known as a "supercage," which is more accurately termed a supercavity, accessible through a three-dimensional 12-ring pore system Extensive studies have been conducted on FAU zeolites exchanged with various cations, including Na+, K+, Ba2+, Cu2+, Ni2+, Zn2+, and Sr2+.
FAU-type zeolite can be represented as
Zeolites are categorized based on their aluminum (Al) content per unit cell, represented as |M 𝑥 (H 2 O) 𝑦 | [Al 𝑥 Si 192 – 𝑥 O 384 ] – FAU, where 'M' denotes a monovalent cation or half of a divalent cation The number of Al atoms can range from fewer than 4 to 96, resulting in Si/Al ratios from 1 to over 50 Specifically, zeolite X contains between 77 and 96 Al atoms per cell, corresponding to Si/Al ratios of 1 to 1.5, while zeolite Y has fewer than 76 Al atoms per cell, with Si/Al ratios exceeding 1.5.
The FAU zeolite framework is characterized by the interconnection of sodalite cages via double six-rings, creating a complex three-dimensional channel system This system features 12-ring pores that lead into expansive supercages Additionally, the positions of Na+ cations within the FAU structure are noteworthy.
1.1.1 Characteristics of the NaX zeolite
Numerous studies have explored the cation sites within the FAU framework, which typically feature cations positioned at the diagonals of cubic unit cells, designated as I, I′, II, II′, III, III′, and U Specifically, site I is located at the center of the hexagonal column cage, while site I′ is situated 1 Å away from this center within the β cage (sodalite cage) Site II is found in the supercage on the face of the six-ring, with site II′ in the sodalite cage linked to the supercage Additionally, sites III and III′ are near the supercage wall, where a cation at site III can easily transition to site III′ with minimal perturbation Lastly, site U is positioned at the center of the β cage, making these cation sites critical for interactions with adsorbed molecules.
II and III, as the others are buried in cages accessible only through six-rings
Recent research by Frising and Leflaive highlights the extensive examination of NaX zeolites, which serve as precursors for nearly all exchanged X zeolites Key findings from these studies are summarized in Table 1.1, showcasing the diverse analytical methods used to determine the structures of the tested samples.
Recent studies utilizing techniques such as X-ray diffraction (XRD), nuclear magnetic resonance spectroscopy, and neutron diffraction have revealed the structural composition of the FAU unit cell, which contains 16 D4R, eight sodalite cages, and eight supercages The cation distribution within dehydrated NaX zeolites shows a preference for sites that enhance interaction with framework oxygen atoms while minimizing cation–cation repulsion Initially, sites I and II accommodate up to 48 cations, with site I′ becoming favorable as the cation count exceeds this threshold, allowing for up to 64 cations When more than 64 cations are present, excess cations occupy site III in the supercage The introduction of water or other adsorbates can alter this cation distribution, as seen when Na cations shift from site I to site I′ to optimize interactions with both framework oxygens and water molecules Notable findings from Olson and Takaishi will be discussed in detail.
Olson utilized single-crystal X-ray diffraction (XRD) to identify the positions of sodium ions in dehydrated NaX zeolite, characterized by the formula Na88Al88Si104O384 and a cubic structure (Fd3) with a lattice parameter of 25.009 Å In each unit cell, the sodium ion occupancy was found to be 2.9 at site I, 21.1 at site I′, and 31.0 at another designated site.
Takaishi's model effectively elucidates the cation distribution within the FAU structure, demonstrating strong alignment between theoretical calculations and experimental findings The model identifies D6Rs containing either three or six Al atoms, which exhibit three-fold symmetry and serve as building blocks for the tertiary building unit (TBU) that comprises four D6Rs and 48 T sites The FAU unit cell is composed of four TBUs, totaling 16 D6Rs and 192 T sites Notably, a tetrahedral arrangement of four TBUs forms a β-cage with 24 T sites at its center, enabling the construction of a unit cell featuring eight β-cages and 192 atoms The Na88X configuration includes three types of cages in specific quantities, with the distribution of Na+ ions within the sodalite cage depicted Additionally, various factors such as cage size, static electric field distribution, cation radius, and zeolite hydration states further influence cation distribution.
The illustration in Figure 1.2 depicts TBUs containing either 24 or 21 aluminum (Al) atoms, with a SOD cage positioned at the center The Al atoms are represented by the (•) symbol The configurations include TBU (24Al) and two variations of TBU (21Al) Additionally, the figure shows the connection of four TBUs resulting in a newly formed SOD cage, indicated by dotted lines Furthermore, the arrangement of sodium (Na+) ions within the SOD cages is illustrated, highlighting three distinct cages: cage 1, cage 2, and cage 3.
The distribution of cation sites in NaX zeolites is influenced by various experimental conditions and methods This study examines multiple types of cation sites, focusing on their hydration states An analysis method is employed to assess these factors, providing valuable insights into the characteristics of NaX zeolites.
9.4 H y d ra ted S ing le c ry stal XRD XRD a cc ompl ished a t ro om t emper ature
R eh y dr ated S ing le c ry stal XRD XRD a cc ompl ished a t ro om t emper ature
XRD a cc ompl ished a t ro om t emper ature
O 3 8 4 9 20 31 23 H y d ra ted S ing le c ry stal XRD [40 ]
O 3 8 4 2.9 29.1 3 1 29.8 400 o C a t 2 x 10 -6 torr S ing le c ry stal XRD [6]
O 3 8 4 8.43 15.68 19.28 12.83 20.16 II * : 1 1.06 Eva cua ted at 523 K P owde r X R D 21 H 2 O/u.c (b y TP D)
1 0 9 6 O 384 29.1 32.3 26.9 Eva cua ted d eh y dr ated P owde r ne utron dif fr ac ti on C he mi sorbe d meth y li um ions [ 44 ]
Zeolite modification, primarily through secondary synthesis methods, enhances specific properties for targeted applications, including acidity, thermal and hydrothermal stability, and catalytic performance This process improves channel structures, surface properties, microporous frameworks, and charge-balancing ions Additionally, these modification techniques can create novel properties unattainable by direct synthesis.
The properties and functions of zeolites are significantly influenced by their framework, pore structure, and the nature and number of charge-balancing cations These cations play a crucial role in determining zeolite characteristics, with multivalent cations enhancing the stability of the zeolite structure under thermal and hydrothermal conditions As a result, extensive research has been conducted on the cation-exchange capabilities of zeolites, leading to comprehensive documentation of their ion-exchange properties.
Zeolite modification has primarily relied on four methods: cation-exchange, dealumination, isomorphous substitution of heteroatoms, and channel and surface modification This article summarizes key findings related to the cation-exchange modification of NaX zeolites.
Bendenia et al demonstrated that the crystallinity, texture, and properties of NaX zeolites are significantly influenced by the substitution of Na+ ions with Ni2+ and/or Cr3+ ions Their findings revealed that the loss of crystallinity and micropore volume was more pronounced in Cr(x)X and CrNi(x)X zeolites compared to Ni(x)X zeolites Additionally, the interactions between carbon dioxide and Cr(x)X zeolites were found to be stronger than those with Ni(x)X zeolites The study also delved into the acidity of these ion-exchanged zeolites.
ray spectroscopy (EDS)
Nitrogen adsorption isotherms
Zeolites possess micro-, meso-, and macro-sized pores that significantly influence their catalytic and separation properties According to IUPAC terminology, micropores are defined as having a diameter smaller than 2 nm, while mesopores range between 2 nm and 50 nm.
Nitrogen adsorption analysis is essential for determining the pore size and surface area of zeolites, with micropores measuring 2 to 50 nm and macropores exceeding 50 nm The characteristics and configuration of these pores can be inferred from the nitrogen adsorption and desorption isotherms As per IUPAC recommendations, adsorption isotherms are categorized into six distinct types (I–VI), illustrated in Figure 1.4.
Type I physisorption adsorption isotherms, as defined by IUPAC in 1985, are characterized by a horizontal plateau that indicates an asymptotic value for the mass adsorbed, which remains constant even under high gas pressures These isotherms are typically associated with microporous materials, where adsorption occurs through micropore filling without the formation of multilayers.
Type II: These isotherms typically describe adsorption in mesoporous materials that exhibit monolayers at low pressures and disperse, nonporous, or exclusively macroporous solids (pore diameters > 50 nm) at
23 higher pressures near the saturation of multilayer adsorption These isotherms can often be described by the Brunauer–Emmett–Teller (BET) equation or its generalizations (Eq 1.4)
Type III: These isotherms are found for systems where the adsorbate–sorbent interaction is weaker than the adsorbate–adsorbate interaction, i.e., strongly associating admolecules
Type IV: These isotherms describe the adsorption behavior of special mesoporous materials showing pore condensation together with hysteresis behavior between the adsorption and desorption branches
Type V: These isotherms are similar to type IV curves but with nearly perpendicular middle portions of the adsorption and desorption branches, often near fairly high gas pressures These portions reflect the presence of mesopores in which phase changes, such as pore condensation, may occur
Type VI: These isotherms represent stepwise multilayer adsorbates; the layers become more pronounced at low temperatures
The presence of a hysteresis loop in isotherms signifies the existence of mesopores, with the loop's shape reflecting their geometric characteristics The IUPAC classification identifies four distinct types of hysteresis loops.
Figure 1.5 Types of hysteresis loop of adsorption isotherms (IUPAC 1985)
H1: reflects a system described by an independent pore model In principle, the pore size can be determined from both the desorption and adsorption branches if appropriate methods are available
H2: caused by pore blockage/percolation or cavitation phenomena in
25 mesoporous and micro/mesoporous pore networks
Pore blocking: Pore (cavity) size can be obtained from the adsorption branch; neck size can be obtained from the desorption branch
Cavitation : Pore (cavity) size can be obtained from the adsorption branch; no pore size information can be obtained from the desorption branch
H3/H4: observed in very disordered micro/mesoporous pore networks and caused by a combination of various phenomena (e.g., cavitation, pore blocking) Pore (cavity) size can be obtained from the adsorption branch
Several methods are available for the analysis of nitrogen physisorption data, such as the Langmuir isotherm, BET, t-plot, and Barrett– Joyner–Halenda (BJH) methods
Langmuir adsorption isotherm [56]: Under linearization, the
Langmuir equation for type I isotherms, 𝑊
𝑚 (Eq 1.2) where P is the adsorbate pressure, K is the Langmuir constant, and 𝑊 𝑊⁄ 𝑚 is the weight adsorbed in the system relative to the weight adsorbed in a
A plot of \( P_W/P \) results in a straight line with a slope of \( 1 \, W/m \) and an intercept of \( 1 \, KW/m \), allowing for the calculation of both \( K \) and \( W_m \) Once \( W_m \) is determined, the sample surface area \( S_t \) can be calculated using Equation 1.3.
𝑀 ̅ (Eq 1.3) where 𝐴 𝑥 and 𝑀̅ are the cross-sectional area and the molecular weight of the adsorbate, respectively, and 𝑁̅ is Avogadro’s number
Brunauer, Emmett, and Teller (BET) theory [57]: In 1938,
Brunauer, Emmett, and Teller expanded Langmuir’s kinetic theory to address multilayer adsorption, introducing the BET theory This theory posits that the top molecules in adsorbed layers exist in dynamic equilibrium with the vapor When only a single layer of adsorbate covers the surface, equilibrium is established between this layer and the vapor; similarly, with two layers, the upper layer remains in equilibrium with the vapor, and so on Although the specific locations of surface sites covered by one or more layers may fluctuate, the number of molecules in each layer remains constant The BET equation encapsulates this relationship.
27 where C is a constant, P/P0 is the relative pressure of the adsorbate, W is the weight adsorbed, and Wm is the weight adsorbed in a completed monolayer
In the region of relative pressure near a completed monolayer (0.05
< 𝑃 𝑃⁄ 0 < 0.3), a plot of 1 𝑊⁄ [𝑃 𝑃⁄ 0 − 1] versus 𝑃 𝑃⁄ 0 will give a straight line of slope (𝐶 − 1) 𝑊⁄ 𝑚 𝐶 and intercept 1 𝐶𝑊⁄ 𝑚 , from which both C and Wm can be calculated
The total surface area can be calculated as follows:
The specific surface area can be calculated using the equation \( S_t = W_m \overline{M} N_A x \), where \( A_x \) represents the cross-sectional area and \( \overline{M} \) denotes the molecular weight of the adsorbate To find the specific surface area, simply divide \( S_t \) by the weight of the sample.
The BET method is widely used for analyzing nitrogen physisorption data; however, its calculated surface area lacks physical significance for microporous materials due to unmet boundary conditions for multilayer adsorption Instead, the values derived from this method should be viewed as proportional to the total micropore volume rather than reflecting specific surface area Conversely, the t-plot method effectively calculates both micropore volume and mesopore area, allowing for the determination of zeolite quantities based on the micropore volume.
The t-plot method [58, 59]: Originally proposed by Lippens and de
The t-plot method, developed by Boer, compares the isotherm of a microporous material with a standard type II isotherm, allowing for the determination of micropore volume and surface area, as well as insights into average pore size This method employs a composite t-standard curve derived from nonporous adsorbents with similar BET C constants to the microporous sample The experimental isotherm is then transformed into a t-curve, plotting the volume of gas adsorbed against the standard multilayer thickness (t) on the reference non-porous material at the corresponding P/P0 The t values are calculated using a thickness equation that describes the specific standard curve.
One popular thickness equation is that obtained by de Boer, which represents nitrogen sorption at 77 K on nonporous adsorbents with oxidic surfaces, such as siliceous materials:
Figure 1.6 (A) Standard type II isotherm; (B) t-plot from type II isotherm;
(C) type II isotherm + microporous sample; (D) t-plot from isotherm C; € isotherm of a microporous material; (F) t-plot from E; (G) Type IV isothenn; and (H) t-plot from type IV isotherm
The experimental isotherm and the standard isotherm exhibit distinct shapes in the non-linear regions of the t-plot, revealing positive or negative intercepts when extrapolated to t = 0 These deviations from the standard isotherm provide valuable insights into the micropore volume and micropore surface area of the adsorbent.
Figure 1.6 illustrates typical t-plots for both microporous and non-microporous samples When isotherm A resembles the standard isotherm of a nonporous sample (type II), t-plot B appears as a straight line that intersects the origin, with its slope representing the surface area, as defined by equation (Eq 1.7) By applying the slope, s, from the plot in Fig 1.6B, equation (Eq 1.7) simplifies to equation (Eq 1.8).
𝑆 𝑡 × 10 4 (Eq 1.7) where 𝑡 is the statistical thickness and 𝑆 𝑡 is the total surface area, which was calculated from nitrogen adsorption at 77 K
In the absence of micropores, there is a strong correlation between the t-area (𝑆𝑡) and the surface area measured by the BET method However, when a small fraction of micropores is present, the adsorption isotherm C indicates a higher gas uptake at low relative pressures, while the t-plot D reflects this change.
31 linear and, when extrapolated to the adsorption axis, will show a positive intercept, i, equivalent to the micropore volume 𝑉 𝑚𝑖𝑐𝑟𝑜 :
The external surface area (𝑆 𝑒𝑥𝑡) of a microporous sample can be determined from the slope of the t-plot, denoted as D Consequently, the micropore surface area (𝑆 𝑚𝑖𝑐𝑟𝑜) can be calculated using the equation 𝑆 𝑚𝑖𝑐𝑟𝑜 = 𝑆 𝐵𝐸𝑇 − 𝑉 𝑚𝑖𝑐𝑟𝑜 (𝑐𝑚 3 ) = 𝑖 × 0.001547 This relationship highlights the importance of accurate measurements in assessing the surface characteristics of microporous materials.
The slope of the straight line in Fig 1.6D reflects the external surface area due to the complete filling of micropores The adsorption isotherm E is characteristic of samples containing solely micropores Similarly, the t-plot F is analyzed in the same way as D In contrast, a typical type IV isotherm, indicating the existence of mesopores, is illustrated in G This type of isotherm would correspond to the t-plot shown in H, assuming the absence of micropores.
O VERVIEW IN H YDROGEN S ULFIDE A DSORPTION
Hydrogen sulfide, commonly known as sewer gas or stink damp, is a highly toxic and flammable gas characterized by its colorless appearance and distinct odor of rotten eggs at concentrations below 1 ppm However, relying on its smell to gauge H2S levels is dangerous, as higher concentrations quickly impair human olfactory senses The detection threshold for H2S ranges from 0.001 to 0.13 ppm, but at 50 ppm, individuals may lose their ability to detect the gas, and at levels exceeding 100 ppm, there is a rapid decline in the sense of smell.
Hydrogen sulfide (H2S) is a crucial raw material for producing sulfur and thioorganic compounds, while also serving as an environmental contaminant It naturally occurs in various sources, including crude petroleum, natural gas, volcanic gases, salt mines, hot sulfur springs, and stagnant bodies of water Additionally, H2S is generated through the biological degradation of organic materials, particularly in anaerobic conditions from human and animal waste, making it prevalent in livestock barns and sewage treatment facilities.
Hydrogen sulfide (H2S) is commonly produced as a by-product in various industrial processes, including petroleum refineries, petrochemical facilities, natural gas processing plants, Kraft paper mills, iron smelting operations, coke production, food processing industries, and tanneries.
[18, 63] Hydrogen sulfide has been used for many important industrial processes as follows [18, 63–65]:
- The production of thioorganic compounds, sodium sulfide, and sodium hydrosulfide
- The purification of hydrochloric and sulfuric acids the replenishment of sulfide content in Kraft pulping
- The purification of ores by selective flotation
- The precipitation of copper, nickel, and cobalt sulfides from ores in metallurgy and the removal of copper, cadmium, and titanium from spent catalysts
- The activation of catalysts, such as presulfiding petroleum cracking catalysts, and catalyst poisoning
- The production of extreme-pressure lubricants
- The formulation of rare earth phosphors for use in color television tubes and the treatment of metal surfaces
- The passivation of the walls of reactors operating at high temperatures in petroleum operations
- The formation of a layer of sulfide on the surface of steel wires or plates that are to be coated with paint or plastic
- The production of heavy water for the nuclear industry and various applications in analytical chemistry
- Finally, it is also used as an additive in cutting oils
General Properties of Hydrogen Sulfide
Hydrogen sulfide (H2S) possesses distinct physical and chemical properties, as detailed in Table 1.3 This gas readily dissolves in water and oil, with the potential for release during heating, depressurization, or agitation of these liquids Due to its greater density than air, H2S can accumulate in low-lying areas, posing significant risks in environments where it may be present Additionally, hydrogen sulfide is highly flammable and explosive, further increasing safety concerns during handling and storage.
H2S gives off sulfur dioxide, another dangerous gas that is toxic, strong- smelling, and irritating to the respiratory system
Table 1.3 The physical parameters of Hydrogen sulfide H2S [18, 66]
Solubility in Water Slightly soluble (398 mg/100 g at 20 o C)
Very soluble in alkanolamines and soluble in the various organic mediums as alcohol, acetone, ether … pH value 4.1 (0.1 N aqueous solution)
Anhydrous hydrogen sulfide gas is practically non-acidic
Water solutions are weakly acidic
Acid Properties: Hydrogen sulfide in aqueous solution is a weak acid with two dissociation constants:
H2S can precipitate the salts of heavy metals (Hg, Pb, Ag, Cu, Cd,
Hydrogen sulfide (H2S) can be removed from natural gas through various processes that involve its absorption by amines This compound, which can exist as sulfides like Co and Ni, dissolves and reacts with amines, and is released upon heating.
Reducing Properties: The standard redox potentials of H2S at 25 °C are as follows:
Hydrogen sulfide undergoes reactions with various oxidizing agents, resulting in sulfur compounds that exhibit different oxidation states The specific products formed, including sulfur oxides (SO2, SO3), salts (S2O3²⁻, SO3²⁻, SO4²⁻), polythionic compounds, elemental sulfur, and sulfanes, depend on the operating conditions and the type of oxidizing species involved.
Hydrogen sulfide (H2S) is produced through the chemical reaction of sulfur with hydrogen or hydrocarbons, as well as by hydrogen reduction and the acid decomposition of sulfides Additionally, H2S can be recovered from natural gas, associated gases, and refinery gases.
Hydrogen sulfide (H2S) is a highly toxic gas that can cause severe irritation to the eyes, nose, throat, and lungs Excessive exposure can inhibit the brain's breathing center, potentially leading to fatal outcomes Immediate first aid is crucial for the possibility of revival This gas poses a significant inhalation risk, particularly at ambient temperatures.
Numerous studies have highlighted the significance of human exposure to H2S; however, the specific exposure levels and durations are often unknown or only roughly estimated The effects of hydrogen sulfide at different exposure levels are believed to vary accordingly.
Table 1.4 The effects of hydrogen sulfide in humans [61, 67 - 69]
Hydrogen sulfide concentration / ppm Effects
1 – 5 Moderately offensive odour, possibly with nausea, or headaches with prolonged exposure
20 -50 Nose, throat and lung irritation, digestive upset
39 and loss of appetite, sense of smell starts to become “fatigued”, odour cannot be relied upon as a warning of exposure
100 – 200 Severe nose, throat and lung irritation, ability to smell odour completely disappears
Potentially fatal build up of fluid in the lungs (pulmonary edema) in the absence of central nervous system effects (headache, nausea, dizziness), especially if exposure is prolonged
Severe lung irritation, excitement, headache, dizziness, staggering, sudden collapse
(“knockdown”), unconsciousness and death within 4 8 hours, loss of memory for period of exposure
500 – 1000 Respiratory paralysis, irregular heart beat, collapse, and death
It is important to note that the symptoms of pulmonary edema, such as chest pain and shortness of breath, can be delayed for up to 48 hours after exposure
Prolonged exposure to low concentrations of 50–100 ppm can lead to symptoms such as a runny nose, cough, hoarseness, and shortness of breath, while higher concentrations may result in bronchitis.
40 pneumonia, and a potentially fatal build up of fluid in the lungs [67, 68]
Reports indicate a concerning number of worker fatalities, particularly in the petroleum, sewage treatment, and agricultural sectors Many of these deaths occur in confined spaces such as sewers, sludge tanks, and pits, where multiple fatalities can arise when would-be rescuers enter these hazardous areas without proper respiratory protection or safety measures Unfortunately, these rescuers often succumb to H2S exposure as well, exacerbating the tragedy.
Workers exposed to serious short-term hydrogen sulfide (H2S) may fully recover or face long-term health effects H2S can impact the nervous system and respiratory functions, with studies indicating potential permanent issues such as fatigue, anxiety, irritability, cognitive decline, reduced attention span, impaired learning and memory, altered smell, and motor deficits These neurological effects may stem from oxygen deprivation to brain cells during severe exposure Additionally, respiratory symptoms linked to H2S exposure can include shortness of breath, chest tightness, and wheezing, reflecting airway hypersensitivity.
41 dysfunction) [72]; permanent lung damage (pulmonary fibrosis) [73]; and significant reductions in residual volume (a measure of lung function) [74]
Effective detection of hydrogen sulfide in the air is crucial, particularly in and around storage areas Permanent detection tools are essential for monitoring this hazardous gas, and sensors must be equipped to trigger both audible and visual alarms when concentrations exceed safe levels.
When hydrogen sulfide (H2S) levels in the air exceed 10 ppm, several sensor models can accurately detect it Electrochemical sensors operate on the principle of gas diffusion, where H2S reacts with an electrolyte in the presence of air or nitrogen, as seen in devices like Honeywell analytical detectors and Sulfytron 4010 Additionally, semiconductor sensors detect H2S by allowing it to diffuse into a specially designed semiconductor, altering its electrical resistance These detectors are effective for measuring hydrogen sulfide concentrations in air or gas mixtures A simple and reliable leak detection method involves lead acetate paper, which turns black upon exposure to H2S, although it does not provide concentration levels.
Recent research has focused on the adsorption and removal of hydrogen sulfide from biogas and gas mixtures, investigating various materials including metal hydroxides, activated carbon, graphene, and different types of zeolites Table 1.5 summarizes the findings from prior studies related to this subject.
Table 1.5 The H2S adsorption capacities of the various absorbents and conditions of test methods
Authors (year) Type of adsorbents
Conditions and test methods Ref
AgY Cu(I)Y Cu(II)Y Cu/AC
By standard gravimetric method at
NaX AgX AgY CuX CuY
10 ppm H2S in He 10/100 ppm H2S in N2
By performing H2S breakthrough experiments at room temperature
Activated carbon, A A_100-NaOH A_100-KOH A_100_KI A_100-Na2CO3
By performing H2S breakthrough experiments at 30 o C/
Zinc-based metal- organic framework, IRMOF-3
By performing H2S breakthrough experiments at room temperature
Fine rubber particle media, FRPM
By performing H2S breakthrough experiments at room temperature
13X 13X-Im-Cu 13X-Im-Zn 13X-Ex-Cu 13X-Ex-Zn
By performing H2S breakthrough experiments at room temperature
Hydrogen sulfide was determined by pH
AC AC-NaOH AC-KOH AC-Na2CO3
1.46 1.69 1.67 2.46 meter based on its acid property
EXPERIMENTAL
2.1 Modification and characterization of the NaX zeolites
2.1.1 Cobalt (II)-exchange processes and thermal treatment
Ionic exchange was performed at room temperature by stirring 5 g of zeolite NaX, spherical tablets with a diameter of approximately 3.0 - 3.5 mm, in 250 ml of Co(NO3)2 aqueous solution The samples underwent thorough washing to remove nitrates and were subsequently dried at 120 °C for 12 hours Co2+ concentration was measured using an atomic absorption spectrometer with a cobalt cathode lamp and air-acetylene flame Zeolite NaX samples exchanged sodium cations for cobalt cations from Co(NO3)2 solutions at various concentrations (0.05 to 0.20 mol L−1), with the pH of the solutions adjusted between 5.00 and 7.00 to assess the influence of cobalt concentration and pH on the ion-exchange process.
Dried samples underwent thermal treatment at varying temperatures of 400, 500, 600, and 700 °C for 3 hours, with a heating rate of 20 °C per minute Following the treatment, the samples were taken out of the furnace and allowed to cool to room temperature Before utilization, the thermally treated samples were prepared accordingly.
The particles were securely stored in air-tight vials to avoid moisture absorption Their structural characterization, size, and morphology were analyzed using X-ray diffraction (XRD) techniques Additionally, nitrogen adsorption isotherms were measured at 77.5 K with a Micromeritics ASAP 2020 instrument.
One hundred grams of zeolite NaX were exchanged with cobalt cations in a glass column (4 mm ID, 40 cm length) using a continuous flow of 2 L of Co(NO3)2 solution at a concentration of pH 6.23 and a flow rate of 1 mL/min The product was washed to remove nitrates and then dried at 120 °C for 12 hours, as the removal of Co2+ and Na+ cations is essential to prevent the formation of NaNO3 and cobalt oxides, which can clog the pore system of NaX zeolite and reduce adsorption capacities The dried samples underwent thermal treatment at varying temperatures (400, 500, 600, and 700 °C) for 3 hours at a rate of 20 °C/min After cooling to room temperature, the thermally treated samples were stored in airtight vials to avoid moisture absorption, and were labeled as n-CoX-T.
47 where n (mol L -1 ) represents the concentration of the Co(NO3)2 solution and
T represents the heating temperature The n-CoX-0 is the sample that was left untreated
The Freundlich and Langmuir isotherm models were employed to analyze the sorption equilibrium data for cobalt (Co) on NaX zeolite Coefficients for these models were determined through linear least-squares fitting.
The Freundlich isotherm equation C Z k F C e 1 / n can be written in linear form as below: e F
The equilibrium concentrations of cobalt (Co) in zeolite and liquid phases are represented by C z (mmol g –1) and C e (mmol L –1), respectively, as described by the equation C = k F log C e + log n (Eq 2.1) The Freundlich constants, k F and n, indicate the sorption capacity and intensity, and can be determined from the slope and intercept of the linear plot.
Under linearization, the Langmuir sorption isotherm equation
The Langmuir constants, Q m and k L, represent the sorption capacity and energy of sorption, respectively These constants can be determined by analyzing the intercept and slope from the linear plot of 𝐶 𝑒 ⁄𝐶 𝑧 against 𝐶 𝑒.
The pseudo-second-order model was used to analyze the sorption kinetic data for cobalt on the various adsorbents [80] The pseudo-second- order equation can be written as follows:
(g mg –1 min –1 ) is the rate constant and q t and q e (mg g –1 ) are the amounts of sorption at time t (min) and at equilibrium, respectively
The integration of (Eq 2.3) and application of the above conditions gives t q k q q e t e 2
After rearranging (Eq 2.4) into a linear form, we obtain q t q q k t e e t
49 k 2 and q e can be obtained from the intercept and slope, respectively, of the plot t q t vs t
The crystallinity and mineralogical composition of Co–X samples were analyzed using X-ray diffraction (XRD) with a PANalytical X’pert PRO system at the National Kaohsiung University of Applied Sciences, Taiwan The analysis utilized CuKα1 radiation over a 2θ range of 5 to 65°, with a step size of 0.02° This technique allows for the assessment of structural changes in the samples after thermal treatment, with the crystallinity percentages of cobalt (II)-treated zeolites calculated from the heights of ten major diffraction peaks, including (1 1 1) and (2 2 0).
The diffraction patterns of the samples revealed peaks at specific 2θ angles: 6.1°, 10.0°, 11.7°, 15.4°, 20.1°, 23.3°, 26.6°, 30.9°, 31.9°, and 33.6°, corresponding to the indices (3 1 1), (3 3 1), (4 4 0), (5 3 3), (6 4 2), (5 5 5), (8 4 0), and (6 6 4) These peaks were analyzed using the Fd-3m cubic space group (227 origins at -3m) Notably, a crystallinity value of 100% was attributed to the NaX zeolite, indicating its maximum degree of crystallinity.
Figure 2.1 The PANalytical X’pert PRO system
SEM imaging was conducted with a JOEL microscope (model JSM
At National Sun Yat-sen University in Taiwan, the technique employed at 6330 F) effectively demonstrates the morphological changes in NaX zeolite and cobalt (II)-exchanged NaX zeolite during cobalt (II) exchange and thermal treatment Additionally, the amount of cobalt uptake was quantified using atomic absorption spectroscopy (AAS) with a cobalt cathode lamp and air acetylene flame.
Figure 2.2 The JOEL microscope (model JSM 6330 F) and Energy dispersive X-ray spectroscopy system
The adsorption–desorption isotherms were collected using an ASAP
In 2020, an instrument from Micromeritics was utilized at the National Kaohsiung University of Applied Sciences to conduct analyses at a temperature of 77.5 K The collected data were evaluated using the BET model for specific surface area and the BJH model to determine pore diameter, pore distribution, and the volumes of micro- and mesopores.
Figure 2.3 The ASAP 2020 instrument (Micromeritics)
Figure 2.4 Experimental system for H2S adsorption
Hydrogen sulfide (H2S) adsorption was performed at room temperature using a specialized adsorption system Before the process, zeolite adsorbents were dried at 100 °C for 12 hours to eliminate moisture The dried zeolite was then placed in a test tank for H2S adsorption, with nitrogen (N2) flow introduced initially to purge the system of air.
The H2S flow was directed through a buffer tank to regulate its concentration using N2 flow Once the H2S concentration stabilized at 100 ppm after closing the shut-off valve (SV3), the flow was redirected through the adsorbent by opening valve SV4 A Honeywell analytical detector monitored the H2S concentration and adsorption duration throughout the process.
The rate of decrease of the hydrogen sulfide concentration was calculated using Eq 2.6: tn tn
(Eq 2.6) where C n 1 , and C n (ppm) are the concentrations of H 2 S at time t n 1 , and t n
It was deemed that the equilibrium state had been reached when the
H2S concentration ceased to change, i.e., the reduction rate, r, was less than 0.02 The associated concentration is the equilibrium concentration of H2S
The H2S adsorption capacity of the adsorbents, q (mg g –1 ), is determined as follows:
The buffer tank volume (V1) and test tank volume (V2) are essential for determining the concentrations of hydrogen sulfide (H2S) in a solution The initial concentration (C0) and equilibrium concentration (C) of H2S, measured in parts per million (ppm), are critical for evaluating the effectiveness of the adsorbent Additionally, the molar mass of H2S (M) and the mass of the adsorbent sample (madsorbent) play significant roles in this analysis.
In this study, nitrogen gas (N2) was utilized due to its inert nature and lighter weight compared to hydrogen sulfide (H2S), facilitating the desorption process during H2S adsorption As N2 constitutes approximately 80% of the atmosphere, using it as a medium for H2S gas removal serves as an effective proxy for atmospheric H2S removal Additionally, N2 is widely available and commercially accessible, making it a practical choice for this application.
In this study, H2S adsorption was assumed to obey first-order kinetics, which can be expressed as adsC dt k
(Eq 2.8) where k ads (min –1 ) is the rate constant and Co and C (ppm) are the H2S initial concentration and at time t (min), respectively
The integrated form of Eq 2.8 is t lnC lnC 0 k ads (Eq 2.9)
Nitric acid (HNO 3 ): 1.5129 g cm −3 of density; 65% purity, SHOWA Chemical Co., Ltd
NaX zeolites: is the spherical tablets with diameter, ~3.0 - 3.5 mm; 99 % purity, Foever Applied Equipment Co., Ltd
Cobalt (II) nitrate hexahydrate , 𝐂𝐨(𝐍𝐎 𝟑 ) 𝟐 𝟔𝐇 𝟐 𝐎: white, solid powder;
99 % purity, SHOWA Chemical Co., Ltd
Hydrogen sulfide (H 2 S) : gas, 95% H2S and 5% N2, Ming Yang Special Gas Co., Ltd
Nitrogen (N 2 ) : gas, 99.99 % purity, Ming Yang Special Gas Co., Ltd
Distilled water (DIW) : molecular formular as H2O, purity > 98%
A JOEL microscope (model JSM 6330 TF)
An atomic absorption spectrometer with a cobalt cathode lamp and air acetylene flame (AAS)
A Honeywell analytical detector for determination of H2S
STUDY ON MODIFICATION OF N𝒂X ZEOLITE
Chapter 3: STUDY ON MODIFICATION OF 𝐍𝐚𝐗 ZEOLITE
3.1 The kinetics of cobalt (II)-exchange processes
This section discusses the modification of NaX zeolite, following the methodology outlined in route 1 of section 2.1.1 It examines the kinetics of cobalt (II) exchange processes, as well as the influence of pH, initial concentration, and stirring time on the exchange efficiency The findings and discussions related to these factors are presented in section 3.1.
3.1.1 Effect of pH, initial concentration and stirring time on exchange
The cobalt (II)-exchanged process was examined across various pH levels to assess the influence of pH on Co²⁺ exchange, as illustrated in Figure 3.1 The findings indicate that Co²⁺ adsorption is pH-dependent, with lower levels of cobalt at acidic pH and increased adsorption at higher pH values In the NaX zeolite framework, sodium ions are attracted to specific sites due to Coulombic forces from surrounding aluminum atoms and interactions with oxygen atoms However, the presence of H⁺ ions can disrupt Si–O–Al bonds, negatively impacting the zeolite structure and reducing the availability of sodium ions Consequently, the cobalt (II) exchange capacity of NaX zeolite diminishes at lower pH levels.
The stability of the NaX zeolite structure at higher pH levels contributes to its increased Co 2+ exchange capacity However, when the pH exceeds 6.5, the amount of exchanged cobalt decreases due to a reduction in Co 2+ concentration in the aqueous solution, which coincides with the condensation of cobalt hydroxide Notably, Co(OH)2 condensation begins at a Co 2+ concentration of 0.1 (mol L -1) when the pH reaches 6.05.
Figure 3.1 Effect of pH on the Co 2+ -exchange at room temperature
Data relating to the effect of the initial concentration of solution and the soaking time in the cobalt exchange at room tempetature is shown in
Figure 3.2 and Figure 3.3 The figures indicate that the cobalt adsorbed increased with increasing the initial concentration of solution up to 0.14 mol
The adsorption of cobalt initially increased rapidly during the first five hours, followed by a slower rate of increase as stirring time continued, while L-1 remained nearly unchanged throughout the process.
Figure 3.2 Effect of the initial concentration in the Co 2+ -exchange at room temperature
Figure 3.3 Effect of the soaking time in the Co 2+ -exchange at room temperature
The sorption isotherm of cobalt on zeolite NaX at room temperature indicates an ion-exchange process, as illustrated in Figures 3.4 and 3.5 The exchange data were analyzed using Freundlich and Langmuir isotherm models, with the fitted constants and regression coefficients detailed in Table 3.1 The graphical representation of the isotherm is presented in Figure 3.3.
Figure 3.4 show that the fit is better under the Freundlich model (R 2 = 0.92) than with the Langmuir model (R 2 = 0.89)
Table 3.1: Freundlich, Langmuir and pseudo-second-order model constants
Figure 3.4 Freundlich sorption isotherm of cobalt (II) cations at pH = 6.05
Figure 3.5 Langmuir sorption isotherm of cobalt (II) cations at pH = 6.05
This study performed the following ion-exchange
The kinetics of ion-exchange, which describes the rate at which cobalt is absorbed, is a crucial factor that determines the capacity of the process The reaction involves the exchange of cobalt ions in aqueous solution with sodium ions, highlighting the dynamic nature of ion-exchange mechanisms.
63 ion-exchange The experimental ion exchange kinetic data were fitted to pseudo-second-order rates, as in (Eq 2.5) shown above
Figure 3.6 Pseudo-second-order sorption kinetic model of cobalt (II) on
The linear observation from Figure 3.6 demonstrates that the ion exchange kinetic data for cobalt on zeolite NaX is accurately represented by the pseudo-second-order model, achieving a high regression coefficient of R² = 0.994.
Co 2+ -exchange (𝑘 2 ) and the equilibrated amount of cobalt were 7.2 x10 -
5(g mg −1 min −1 ), and 125 (mg g −1 ), respectively (the initial concentration of cobalt (II) solution is 0.1 (mol L −1 ) This cobalt (II)-exchanged NaX
64 zeolite was labeled as 0.1 − NaCoX , where 0.1 (mol L −1 ) is the concentration of the cobalt (II) solution
3.2 Characteristic of the cobalt (II)-exchanged NaX zeolites
3.2.1 Effect of treatment temperature on the cobalt (II)-exchanged
Cobalt (II) ions were introduced to NaX zeolite using a 0.1 M cobalt (II) nitrate solution, resulting in the formation of 0.1-NaCoX zeolite This cobalt (II)-exchanged NaX zeolite was then subjected to various treatment temperatures The impact of these treatment temperatures on the properties of the cobalt (II)-exchanged NaX zeolite was analyzed and discussed.
3.2.1.1 X-ray Diffraction and SEM analysis
Figure 3.7 X-ray powder diffraction patterns of NaX and 0.1-NaCoX thermally treated zeolites: thermally untreated sample NaCoX-0 and the samples treated at 400 o C for 3 h (0.1-NaCoX-1), 500 o C for 3 h (0.1-NaCoX-
2), 600 o C for 3 h (0.1-NaCoX-3) and 700 o C for 3 h (0.1-NaCoX-4)
The X-ray diffraction patterns of NaX and the thermal effect on Co 2+ - exchanged zeolite NaX were shown in Figure 3.7 It can be seen that the Faujasite framework of NaX zeolite is retained during the cobalt exchange process However, the rapidly decrease in crystallinity was due to the presence of cobalt ions in the NaX zeolite framework, leading to decreased
The framework structure of NaCoX-4 remained stable up to 600 °C but was compromised at temperatures exceeding 700 °C due to amorphorization during thermal treatment This is evidenced by the disappearance of the (220) and (311) peaks at 2θ = 10.03° and 11.7° in the X-ray diffraction pattern Additionally, the (111) peak at 2θ = 6.14° (d = 14.45 Å) shifted to lower d values following thermal treatment The unit cell parameters were measured for cubic symmetry, revealing values of 24.835 Å for the untreated sample, 24.827 Å for the sample treated at 400 °C for 3 hours, and 24.703 Å for the sample subjected to higher temperatures.
600 o C for 3h This finding agrees with the studies of Jeong et al [13] and the unit cell parameters were decreased due to a dehydration
The XRD patterns of 0.1-NaCoX (Figure 3.7) were analyzed in comparison with the findings of Dalia Jonynaite et al [82] and Jeong et al [13], which confirmed the presence of cobalt oxide peaks, notably at 2θ = 42.4.
[200] plane of CoO (ICDD #45-1004); of carnegeite at 2θ = 21.1 degree as a secondary phases
The SEM images of NaX zeolite and thermally treated cobalt (II)-exchange zeolite samples, illustrated in Figure 3.8, reveal that the cobalt (II) exchange process does not significantly alter the morphology of NaX zeolite (Figure 3.8A-C) The particle sizes maintain a nearly spherical shape throughout the process.
The cobalt (II) exchange does not significantly collapse the crystalline structure of NaX zeolite, as evidenced by XRD diffraction analysis Additionally, the shape and surface of the particles remain intact up to 600°C, after which distortion occurs at higher temperatures, as illustrated in Figure 3.8D.
Figure 3.8 The SEM images of NaX and 0.1-NaCoX thermally treated zeolites: NaX sample (A) and the samples treated at 400 o C for 3 h (B), 600 oC for 3 h (C) and 700 o C for 3 h (D)
Figures 3.9 and 3.10 illustrate the nitrogen adsorption/desorption isotherms for NaX and 0.1-NaCoX zeolites at different temperatures The data reveals that micro and mesopore filling occurs at low pressures due to the narrow pore widths and high adsorption potential, characteristic of monolayer-multilayer adsorption typical of type II isotherms Notably, a plateau is not reached at p/po ≥ 0.8, indicating that mesopore adsorption plays a significant role in the overall adsorption process.
Figure 3.9 t-plot isotherms of NaX and thermally treated 0.1-NaCoX zeolites
Figure 3.10 N2 adsorption (solid symbols)/desorption (open symbol) isotherms and pore size distributions (small picture) for NaX and thermally treated cobalt (II)-exchanged zeolites NaX
The adsorption isotherms for mesoporous materials exhibit a type IV isotherm, characterized by a hysteresis loop In the case of 0.1-NaCoX zeolites, the contribution from microporous adsorption is lower compared to NaX zeolites This reduction is attributed to cobalt exchange, which enhances the porosity and roughness of the zeolite surface This effect is evidenced by an increase in the slope of the middle region of the adsorption isotherm curves for NaX zeolites following cobalt (II) exchange.
HYDROGEN SULFIDE ADSORPTION BY MODIFIED
4.1 Characterization of the thermally treated cobalt (II)-exchanged
The study investigated the influence of cobalt (II) on the crystal structure of NaX zeolite, along with the impact of thermal treatment on cobalt (II)-exchanged NaX, utilizing X-ray diffraction (XRD) analysis The findings, including crystallinity percentages and unit cell dimensions, are detailed in Table 4.1.
Table 4.1 The cobalt uptake and the structure and textural properties of cobalt (II)-exchanged zeolites: untreated sample (0.1-CoX-0) and treated at 400, 500, 600, and 700 o C for 3 h
The cobalt (II) exchange process maintains the Faujasite framework of NaX zeolite, although it significantly reduces the percentage of crystallinity due to the incorporation of cobalt (II) ions, which affects the unit cell dimensions During thermal treatment, the 0.1-CoX(s) zeolite retains its framework structure up to 600 °C, but begins to deteriorate at temperatures of 600 °C and above.
At 700 °C, the X-ray diffraction pattern of 0.1-CoX-700 showed the disappearance of the (2 2 0) and (3 1 1) peaks at 2θ = 10.03° and 11.7°, indicating amorphorization The reflection peak positions varied with the treatment temperature, with the (111) peak at 2θ = 6.14° (d = 14.45 Å) shifting to lower d-spacing values after thermal treatment The calculated unit cell parameters for cubic symmetry were 24.835 Å for the untreated sample, 24.827 Å for the sample treated at 400 °C for 3 hours, and 24.703 Å for the sample treated at higher temperatures.
Heating to 600 °C for 3 hours resulted in a decrease in the unit cell parameter, a finding consistent with Jeong et al (2012) This reduction can be attributed to dehydration, as noted by Borissenko et al (2008).
82 increase in the cobalt uptake when the treatment temperature increases (Table 4.1)
Figure 4.1 X-ray powder diffraction patterns of NaX and 0.10-CoX thermally treated zeolites NaX
Figure 4.2 displays SEM images comparing NaX zeolite and thermally treated cobalt (II)-exchanged zeolite samples The analysis reveals that there are no significant morphological differences between the two, as both types exhibit nearly spherical crystallites measuring approximately 1 to 3 micrometers in size This indicates that cobalt (II) exchange does not lead to a substantial collapse of the zeolite structure.
The crystalline structure of NaX zeolite, confirmed by XRD diffraction, reveals that cobalt (II)-exchanged zeolite exhibits dark regions due to a morphous structure resulting from high-temperature thermal treatment While the shape and surface of the particles remain intact up to 600°C, they become distorted at higher temperatures Notably, the cobalt (II)-exchanged zeolite treated at 500°C displays smooth and clear particle surfaces, indicating that the optimal treatment temperature lies within this range.
500 to 600 o C By the way, the morphology of 0.15-CoX-T and 0.2-CoX-T samples were also considered, however, the morphological evolution was similar to 0.1-CoX-T samples
Figure 4.2 SEM images of NaX and 0.10-CoX thermally treated zeolites:
NaX sample (A) and the samples treated at 400 o C for 3 h (B), 500 o C for 3 h (C), 600 o C for 3 h (D) and 700 o C for 3 h (E)
Figure 4.3 N2 adsorption (solid symbols)/desorption (open symbol) isotherms for NaX and thermally treated cobalt (II)-exchanged zeolites NaX
Figure 4.4 Pore size distributions for NaX and thermally treated cobalt
The nitrogen adsorption and desorption isotherms for NaX and 0.1-CoX-T zeolites, subjected to different temperature treatments, are illustrated in Figure 4.3 Based on IUPAC guidelines, these adsorption isotherms are classified as type [insert type here].
IV isotherms illustrate the properties of mesoporous materials, highlighting that micropore and small mesopore filling occurs at low pressures due to high adsorption potential and narrow pore widths Notably, a plateau is not reached in this scenario.
87 p/po ≥ 0.8, obviously indicating expansion of the pore
Table 4.2 The surface characteristic parameters of NaX zeolites and
Cobalt (II)-exchanged NaX zeolites exhibit significant adsorption potential, as indicated by the steeper slope of the adsorption isotherm curves within the p/po range of 0.1 to 0.8 compared to NaX zeolites This suggests that the microporous adsorption in cobalt (II)-exchanged zeolites is lower than that in NaX zeolites, highlighting the impact of cobalt (II) exchange on adsorption characteristics.
88 into a zeolite makes the zeolite surface becomes more porous and rough, here, the micro- and mesoporosity increase
The adsorption isotherms of the 0.1-CoX-T samples indicate that the optimal thermal treatment temperature is between 500 and 600 °C At 700 °C, the zeolites experience surface texture damage, leading to a less porous and rough structure The specific surface areas and micropore surface areas of the thermally treated 0.1-CoX-T samples decrease with higher temperatures due to material compression, while mesopore size and volume significantly increase up to 600 °C The nitrogen adsorption/desorption isotherms reveal a type H3 hysteresis loop, indicative of non-rigid aggregates of plate-like particles and the presence of small mesopores Additionally, capillary condensation in meso- and macropores is evidenced by the tensile strength failure of the nitrogen meniscus.
The analysis of pore size distribution from the adsorption branch reveals a realistic pore size, as illustrated in Figure 4.4 and detailed in Table 4.2 Notably, the mesopore size tends to increase with rising temperature.
Thermal treatment results in a 89% collapse of zeolite texture, significantly increasing mesopore volume Following cobalt (II) exchange, the mesopore size of NaX zeolite decreases markedly, attributed to the differing diameters of the metal cations involved.
4.2 Hydrogen sulfide adsorption by modified coblt (II)-exchanged
This study investigates the adsorption kinetics of H2S, which are modeled as first-order kinetics The rate constant was determined experimentally by plotting lnC against time, as illustrated in Figure 4.5 The values of k ads and R² for the 0.10-CoX-T and 0.15-CoX-T adsorbents are presented in Table 4.3 Notably, all R² values for the experimental data exceeded 70%, indicating a good fit for the kinetic model.
Table 4.3 The H2S adsorption constants of NaX zeolites and Co 2+ - exchanged zeolites
Figure 4.5 The first-order H2S adsorption kinetic model of thermally treated cobalt (II)-exchanged zeolite NaX: (A) 0.1-CoX-T, (B) 0.15-CoX-T
4.2.2 Effect of the treatment temperature on H 2 S adsorption
Hydrogen sulfide adsorption on thermally treated cobalt (II)-exchanged zeolite NaX adsorbents is illustrated in Figure 4.6, with detailed adsorption capacities and times provided in Table 4.3 for zeolites NaX, 0.10-CoX-T, and 0.15-CoX-T at varying temperatures (400, 500, 600, and 700 °C) The thermal treatment significantly enhances the sulfide adsorption capacities of the cobalt (II)-exchanged zeolites, with optimal performance observed at 600 °C, where the adsorption capacities reach 4.241 mg g^-1 for 0.10-CoX-600 and 4.423 mg g^-1 for 0.15-CoX-600 As the treatment temperature increases to 600 °C, both the adsorption capacity and rate constant improve due to increased mesopore size and volume However, at 700 °C, a decrease in adsorption capacity and time occurs, attributed to excessively large mesopore sizes that hinder H2S storage These findings align with the results obtained from XRD analysis and SEM and BET methods.
500 and 0.15-CoX-500 are the highest; however, their sulfide adsorption capacity is less than the cobalt (II)-exchanged zeolites that were treated at
Figure 4.6 H2S adsorption on different zeolite adsorbents: (A) 0.1-CoX-T, (B) 0.15-CoX-T thermally treated cobalt (II)-exchanged zeolite NaX
In addition to physisorption, hydrogen sulfide (H2S) chemisorption occurs on Co2+-exchanged zeolite NaX through an acid-base reaction Research by Jeong et al demonstrated that thermal treatment at elevated temperatures influences this process.