# Comparative Investigation on Two Synthesizing Methods of Zeolites for Removal of Methylene Blue from Aqueous Solution

Authors: Misikir Tamiru Asefa; Gebisa Bekele Feyisa
Journal: International Journal of Chemical Engineering
 (2022)

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## Abstract

Organic dyes discharged from industries have significant effect on ecosystem and health of human being because of their toxicity and appearing colour in the wastewater. Absorption method is a more preferable method than other wastewater treatment methods due to its characteristics of being eco-friendly, simple, and efficient. Zeolites are among the porous materials often used as absorbent of organic dye from wastewater. However, wide use of zeolite has been limited due to its expensive precursors and synthesized methods (i.e., hydrothermal method which needs expensive autoclave). In this work, cheap and widely available precursors aluminum from waste food packaging aluminum foil and low cost silica from sugar cane bagasse ash were used to synthesized zeolite without hydrothermal method (Z-B), where hydrothermally synthesized zeolite (Z-A) was used as a reference. The XRD patterns revealed that Z-B was sodalite octahydrate zeolite and Z-A was zeolite Linde Type A (LTA). The morphology and type of bond in both zeolites were investigated by SEM and FTIR. The synthesized zeolites were used as absorbents for absorbing methylene blue (MB) from aqueous solutions. The MB removal efficiency of the synthesized zeolites was evaluated by using UV-Visible spectroscopy. The results indicate that the absorption capacities of Z-B and Z-A were 3.5 mg/g and 3.9 mg/g at 40 mg/L, respectively. Optimum removal efficiencies of both zeolites were observed at PH of 7 and adsorbent dosage of 0.005 mg/L. The stabilities of both zeolites were tested three times. The absorption isotherms of sodalite octahydrate zeolite and zeolite LTA were effectively fitted with the Freundlich and Langmuir modes. Moreover, the absorption kinetics of both zeolites follow pseudo-second-order kinetics. Therefore, nonhydrothermally synthesized zeolite is alternative absorbent for dye removal due to its safety, cheap cost, using low cost and widely available precursors, and using easy and safe synthesizing method.


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## Body

## 1. Introduction

Water is the most abundant and essential resource for life; life cannot be sustained beyond a few days without water [1, 2]. However, the distribution of water on the Earth’s surface is extremely uneven and also it is contaminated by different types of impurities from different sources such as 2,4-Dinitrophenol (DNP) [3], heavy metals [2], and organic dyes [4]. The discharges from the industries are the most responsible contaminants; for instance, dye effluents are discharged from manufacturing, dyeing, printing, and textile industries [5]. Due to the increase in disposal of wastewater, many methods were proposed for wastewater treatment. The most common methods for dyes removal are chemical precipitation [6], flotation [7], flocculation [8], ion exchange [9], chemical oxidation [10], reverse osmosis [11], ultrafiltration [11], electrodialysis [12], and adsorption [13, 14]. Adsorption is a preferable method compared to other methods due to its easy operation, low cost, flexibility in design, high efficiency, availability, high-quality treated effluent, and its recyclability [4]. Some of the recently reported absorbents for removal of organic dyes from wastewater are carbon-based adsorbent materials [15], sugar cane bagasse [16], rice husk [17], magnetic multiwalled carbon nanotubes-loaded alginate [18], silica supported chitosan/glutaraldehyde [19], pomegranate peels based activated carbon [20], and other organic natural materials [21, 22]. Zeolites are among the best adsorbents because they are composed of three-dimensional crystalline and hydrated aluminosilicates made from the interlinked tetrahedra of silica (SiO4)4+ and alumina (A1O4)5+ [23]; in addition to this, zeolites have big scientific and industrial significance because of their good ion exchange properties, high surface area, nontoxicity, low cost, abundance, and a hydrophilic characteristic [5, 23]. In fact, there are natural zeolites which are less applicaple compared to the synthetic one because of impurities triggered inside the pores [24]. The common method of synthesizing zeolites is the hydrothermal method [25]. However, the hydrothermal method requires expensive autoclave and it is an unsafe method because it needs high pressure and temperature for a long time. In addition to this, everywhere unavailable precursors such as sources of aluminum oxides and silicon dioxides were used to fabricate zeolite [25]. In order to widely use zeolite for many applications, comprehensively available and low-cost raw materials and cheap and effective synthesizing method are substantial. For instance, recently reported lists of raw materials and synthesizing method for fabricating zeolite are reviewed as shown in Table 1. All the mentioned raw materials were used as sources of sodium silicate, but sodium aluminate salt was used as source sodium aluminate which is somehow expensive. In addition to this, the synthesizing method (i.e., the hydrothermal method) requires expensive autoclave which is not available particularly in developing countries. In this work, comparative study of two synthesizing methods such as coprecipitation and hydrothermal methods for synthesizing zeolites which were used for removal of Methylene blues (MB) from aqueous solution was carried out. In addition to this, the bagasse ash was used as source of sodium silicates and food package aluminum foil wastes were turned to the source of sodium aluminate. To our knowledge, no works had been reported on cost free raw materials and comparative study on synthesizing methods of zeolites. The results indicate that the removal efficiency of Z-B was comparable to that of Z-A at PH of 7, adsorbent dosage of 0.005 mg/L, contact time of 150 min, and MB concentration of 10 ppm.Table 1
The lists of raw materials and synthesizing methods of zeolites.
N/sRaw materialsSynthesizing methodsApplicationsReferences1Coal fly ashHydrothermalDesulphurization of wastewater[26]2Rice husk ashHydrothermalCO2[27]3Textile waste ashHydrothermalPb[28]4Bagasse fly ashHydrothermalHeavy metals[29]5KaolinHydrothermal---[27]


## 2. Materials and Methods

### 2.1. Materials

The reagents used in this work were of analytical grade. Hydrochloric acid (35% HCl), sodium hydroxide (NaOH), and metaylene blue (MB) were purchased from Loba Chemical Company. Sugar cane bagasse ash for this experiment was collected from Wonji sugar factory, Ethiopia. All aqueous solutions were prepared using distilled water. Test tubes were used to hold MB and zeolites solution. Waste food package aluminum foils were collected from deposed area in Adama city. Laboratory clock was used to measure time, and oven was used for drying the samples. Filter paper was also used to filter the precipitate part from the solution. Knife was used to cut the aluminum foil into pieces. Magnetic stirrer was used to uniformly mix the solutions.


### 2.2. Methods

#### 2.2.1. Collection and Preparation of Sugar Cane Bagasse Ash

Sugar cane bagasse ash for this experiment was collected from Wonji sugar factory, which is 13 Km far from “Adama” town, Ethiopia. Then, the collected bagasse ash was washed repeatedly with distilled water to remove different impurities, especially salt and organic compounds, sun-dried for 2 days in the laboratory, sieved by 200μm sieves, and stored at room temperature for the next experiment.


#### 2.2.2. Preparation of Sodium Silicate Solution

To obtain pure sodium silicate solution, sugar cane bagasse ash was used as a silica source which is a waste product of sugar factory burned in furnace for 3 hours at 600°C. Calcination can be used to adjust its chemical composition, particularly by reducing its carbon content. Then SCBA was mixed with sodium hydroxide in a ratio of 1 : 1.3, respectively, followed by burning at 600°C for 1 hour. After cooling down the mixture, it was mixed with 200 ml of distilled water and stirred for 10 hours. Then the solution was filtered by using filter paper to obtain sodium silicate solution.


#### 2.2.3. Preparation of Sodium Aluminate

A waste food packaging Al foil was washed and cut into small pieces. 27 g of the powder was dissolved in 250 ml distilled water with 1 M NaOH. Then, the solution was stirred continuously at 60°C with magnetic stirrer until homogeneous solution is obtained [30].


#### 2.2.4. Synthesis of Zeolite

The prepared sodium aluminate solution in the first part by dissolving Al metal and prepared sodium silicate solution was mixed under fast stirring for 10 hours and then left at room temperature to allow the transformation of the mixture from sol into a gel [31, 32]. The gels were transferred to an oven at 100°C for 24 hours and in the section two samples were prepared. For the preparation of the 1st sample, the solution was put in an autoclave and the 2nd sample solution in the beaker and they are abbreviated as Z-A and Z-B, respectively. Then, dry solids were washed repeatedly by distilled water to drop the pH and then dried in the oven. Finally, dried solids were ground into a powder and subsequently calcined at 550°C for 6 hours.


### 2.3. Characterizations

Determination of the crystal structure of phases purity of each sample was characterized by using X-ray powder diffraction at copper Kα radiation (λCuKa = 1.5418 Å), a scan speed of 3.0000 (deg/min), voltage of 40 kV, current of 30 mA, and scanning range of 10–80°. The surface morphology was analyzed by scanning electron microscopy (SEM) and, finally, the chemical bonds and functional groups of the synthesized zeolite and other samples were determined by using FTIR instrument (FT/IR-6600 type A) with a wavenumber range of 200–4000 cm−1. The absorption efficacies of zeolites were evaluated using UV spectroscopy. Langmuir and Freundlich models were utilized to investigate the nature absorption process. Isothermal models such as pseudo-first-order and pseudo-second-order fitting were also used to observe the nature of both zeolites.


### 2.4. Adsorption Experiment

The stock solution was prepared in distilled water by using a standard concentration of methylene blue. A series of adsorption tests were carried out at typical equilibrium conditions for the optimization of adsorbents. The mixture was stirred using a magnetic stirrer during all of the experiments and, finally, samples were analyzed by using UV-Visible spectrophotometer (UV-vis-3600 Plus, Shimadzu). Equations (1) and (2) were used to calculate the removal efficiency and adsorption capacity of the adsorbents: (1)qe=VCo−Cem,(2)η%=Co−CtCo×100,where qe(mg/g) is the equilibrium adsorption capacity, Co is the initial concentration of methylene blue, Ce (mg/L) is the concentration of methylene blue at equilibrium time, Ct is the concentration of methylene blue at a given time, V is the volume of solution, m is mass of the adsorbent, and η is removal efficiency.


## 2.1. Materials

The reagents used in this work were of analytical grade. Hydrochloric acid (35% HCl), sodium hydroxide (NaOH), and metaylene blue (MB) were purchased from Loba Chemical Company. Sugar cane bagasse ash for this experiment was collected from Wonji sugar factory, Ethiopia. All aqueous solutions were prepared using distilled water. Test tubes were used to hold MB and zeolites solution. Waste food package aluminum foils were collected from deposed area in Adama city. Laboratory clock was used to measure time, and oven was used for drying the samples. Filter paper was also used to filter the precipitate part from the solution. Knife was used to cut the aluminum foil into pieces. Magnetic stirrer was used to uniformly mix the solutions.


## 2.2. Methods

### 2.2.1. Collection and Preparation of Sugar Cane Bagasse Ash

Sugar cane bagasse ash for this experiment was collected from Wonji sugar factory, which is 13 Km far from “Adama” town, Ethiopia. Then, the collected bagasse ash was washed repeatedly with distilled water to remove different impurities, especially salt and organic compounds, sun-dried for 2 days in the laboratory, sieved by 200μm sieves, and stored at room temperature for the next experiment.


### 2.2.2. Preparation of Sodium Silicate Solution

To obtain pure sodium silicate solution, sugar cane bagasse ash was used as a silica source which is a waste product of sugar factory burned in furnace for 3 hours at 600°C. Calcination can be used to adjust its chemical composition, particularly by reducing its carbon content. Then SCBA was mixed with sodium hydroxide in a ratio of 1 : 1.3, respectively, followed by burning at 600°C for 1 hour. After cooling down the mixture, it was mixed with 200 ml of distilled water and stirred for 10 hours. Then the solution was filtered by using filter paper to obtain sodium silicate solution.


### 2.2.3. Preparation of Sodium Aluminate

A waste food packaging Al foil was washed and cut into small pieces. 27 g of the powder was dissolved in 250 ml distilled water with 1 M NaOH. Then, the solution was stirred continuously at 60°C with magnetic stirrer until homogeneous solution is obtained [30].


### 2.2.4. Synthesis of Zeolite

The prepared sodium aluminate solution in the first part by dissolving Al metal and prepared sodium silicate solution was mixed under fast stirring for 10 hours and then left at room temperature to allow the transformation of the mixture from sol into a gel [31, 32]. The gels were transferred to an oven at 100°C for 24 hours and in the section two samples were prepared. For the preparation of the 1st sample, the solution was put in an autoclave and the 2nd sample solution in the beaker and they are abbreviated as Z-A and Z-B, respectively. Then, dry solids were washed repeatedly by distilled water to drop the pH and then dried in the oven. Finally, dried solids were ground into a powder and subsequently calcined at 550°C for 6 hours.


## 2.2.1. Collection and Preparation of Sugar Cane Bagasse Ash

Sugar cane bagasse ash for this experiment was collected from Wonji sugar factory, which is 13 Km far from “Adama” town, Ethiopia. Then, the collected bagasse ash was washed repeatedly with distilled water to remove different impurities, especially salt and organic compounds, sun-dried for 2 days in the laboratory, sieved by 200μm sieves, and stored at room temperature for the next experiment.


## 2.2.2. Preparation of Sodium Silicate Solution

To obtain pure sodium silicate solution, sugar cane bagasse ash was used as a silica source which is a waste product of sugar factory burned in furnace for 3 hours at 600°C. Calcination can be used to adjust its chemical composition, particularly by reducing its carbon content. Then SCBA was mixed with sodium hydroxide in a ratio of 1 : 1.3, respectively, followed by burning at 600°C for 1 hour. After cooling down the mixture, it was mixed with 200 ml of distilled water and stirred for 10 hours. Then the solution was filtered by using filter paper to obtain sodium silicate solution.


## 2.2.3. Preparation of Sodium Aluminate

A waste food packaging Al foil was washed and cut into small pieces. 27 g of the powder was dissolved in 250 ml distilled water with 1 M NaOH. Then, the solution was stirred continuously at 60°C with magnetic stirrer until homogeneous solution is obtained [30].


## 2.2.4. Synthesis of Zeolite

The prepared sodium aluminate solution in the first part by dissolving Al metal and prepared sodium silicate solution was mixed under fast stirring for 10 hours and then left at room temperature to allow the transformation of the mixture from sol into a gel [31, 32]. The gels were transferred to an oven at 100°C for 24 hours and in the section two samples were prepared. For the preparation of the 1st sample, the solution was put in an autoclave and the 2nd sample solution in the beaker and they are abbreviated as Z-A and Z-B, respectively. Then, dry solids were washed repeatedly by distilled water to drop the pH and then dried in the oven. Finally, dried solids were ground into a powder and subsequently calcined at 550°C for 6 hours.


## 2.3. Characterizations

Determination of the crystal structure of phases purity of each sample was characterized by using X-ray powder diffraction at copper Kα radiation (λCuKa = 1.5418 Å), a scan speed of 3.0000 (deg/min), voltage of 40 kV, current of 30 mA, and scanning range of 10–80°. The surface morphology was analyzed by scanning electron microscopy (SEM) and, finally, the chemical bonds and functional groups of the synthesized zeolite and other samples were determined by using FTIR instrument (FT/IR-6600 type A) with a wavenumber range of 200–4000 cm−1. The absorption efficacies of zeolites were evaluated using UV spectroscopy. Langmuir and Freundlich models were utilized to investigate the nature absorption process. Isothermal models such as pseudo-first-order and pseudo-second-order fitting were also used to observe the nature of both zeolites.


## 2.4. Adsorption Experiment

The stock solution was prepared in distilled water by using a standard concentration of methylene blue. A series of adsorption tests were carried out at typical equilibrium conditions for the optimization of adsorbents. The mixture was stirred using a magnetic stirrer during all of the experiments and, finally, samples were analyzed by using UV-Visible spectrophotometer (UV-vis-3600 Plus, Shimadzu). Equations (1) and (2) were used to calculate the removal efficiency and adsorption capacity of the adsorbents: (1)qe=VCo−Cem,(2)η%=Co−CtCo×100,where qe(mg/g) is the equilibrium adsorption capacity, Co is the initial concentration of methylene blue, Ce (mg/L) is the concentration of methylene blue at equilibrium time, Ct is the concentration of methylene blue at a given time, V is the volume of solution, m is mass of the adsorbent, and η is removal efficiency.


## 3. Result and Discussion

### 3.1. XRD Analysis

Figure1 shows the XRD results of Z-A and Z-B. In Z-A, the diffraction peaks at angles of 2θ = 13.99°, 19.8°, 24.3°, 31.5°, 34.6°, 37.4°, 42.7°, and 58.09° correspond to the (100), (200), (211), (310), (222), (321), (411), and (440) crystallographic planes, indicating sodalite octahydrate zeolite. These results are aligned with the results reported in [33]. In addition to this, the crystal structure of Z-A was compared with the standard zeolite material and it was found that they have similar structures. The XRD pattern of Z-B was also indicated in Figure 1. Diffraction peaks of Z-B were observed at angles of 2θ = 7.3°, 10.3°, 12.6°, 16.2°, 21.7°, 24.1°, 27.2°, 30°, and 34.2° corresponding to the (100), (001) (101), (110), (101), (101), (210), (001), and (220) crystallographic planes. These data revealed that Z-B is zeolite LTA that is matched with standard of zeolite LTA. We can conclude that the synthesis of zeolite by using different techniques may give different types of zeolite. But, in both cases of zeolites, no considerable amounts of amorphous phases were observed. Here, even though the type of nonhydrothermally synthesized zeolite has different crystal structure as compared with hydrothermal synthesized zeolite, obtaining zeolite without using expensive autoclave is important in order to reduce the cost and for safety issue.Figure 1
XRD patterns of Z-A and Z-B.The crystalline sizes of Z-A and Z-B particles were calculated to be around 23.14 nm and 49.73 nm, respectively, using the Scherrer formula, according to the following equation [34]:(3)D=Kλβcosθ,where λ = 0.15406 nm (wavelength of X-ray), K = 0.9 (Scherrer constant), D is crystalline size (nm), and = FWHM (full width at half maximum). As is expected, the crystal size of hydrothermally synthesized zeolite was smaller than that of nonhydrothermally synthesized one because of the high pressure in the autoclave. Practically, as particle size decreases, the surface area of the material increases, which promotes the absorption of the pollutants in the wastewater [31]. The effect of crystal size on the removal efficiency will be addressed in Section 3.2.


### 3.2. SEM Analysis

Figure2 shows the morphologies of the synthesized zeolites Z-A and Z-B at different magnification. Z-A shows interconnected porous structure with an average particle size less than 24 nm as indicated in Figure 2(a). Similarly, porous structure having interconnection was observed in case Z-B as shown in Figure 2(b). In both cases, spherical-like structures were observed at magnification of 1000x. For the distinct zeolite types, it is reported that different crystal shapes of zeolite are probably associated with different raw materials and preparation methods applied ([32], [33]). Here, although the synthesizing methods were different, the precursors were the same for both zeolites. As a result, a cluster of spherical shapes were observed in both zeolites.Figure 2
SEM image of (a) Z-A and (b) Z-B at different magnification.
(a)(b)


### 3.3. FTIR Analysis

The FTIR spectra of Z-A and and Z-B are shown in Figure3. The FTIR results showed almost similar spectra of packs with different intensity. In the fingerprint region, the spectrum shows a broad and intense band around 1000 cm−1, showing that the characteristic of antisymmetric stretching vibration of the Si-O-Si and a less intense band around 789 cm−1 is due to Si-O-Si symmetric stretching vibrations; this has the same stretching as reported in [35]. Both synthesized zeolites show peaks around 3500 cm−1 and 1650 cm−1, which corresponds to the stretching of the hydroxyl groups (OH-) from adsorbed water molecules. The sharpness of the peaks could determine the amount of water (H2O) molecule adsorbed on the surface of the samples, which in turn can be used to estimate the amount of the hydroxyl functional groups present in the sample as indicated in other research [34]. Other functional groups that exist in the samples are Si-O-Si, indicated by the absorption band at 1020 cm−1, Al-O indicated by peak at 770 cm−1, and Si-O-Al, indicated by the presence of an absorption band at the region between 420 and 494 cm−1. Concerning the existence of functional groups Si, O, and Al observed in the spectrum, it was then concluded that the formation of the zeolite framework was confirmed by the FTIR results as indicated in [36].Figure 3
FTIR spectra of Z-A and Z-B.


### 3.4. Adsorption Performance Analysis

The efficiency of adsorbent in water purification is affected by many factors such as contact time, initial concentration of impurity, adsorbent dose, and pH [37, 38]. In this section, the adsorbent performance is studied with respect to those parameters. Similar to the intrinsic properties of the absorbents, the external factors mentioned have a significant impact on the absorbing properties of absorbents.


#### 3.4.1. Effect of Contact Time on Removal Efficiency of Absorbent

Figure4 shows the effect of the contact time on the MB removal efficiency by both Z-A and Z-B. The adsorption of methylene blue by Z-A and Z-B at room temperature, initial methylene blue concentration of 10 ppm, and adsorbent dosage (Z-A and Z-B) of 3 g/L were used to study the effect of contact time for 150 minutes on the removal efficiency of adsorbents. The adsorptions of methylene blue were slow for both Z-A and Z-B with respect to time because of the low capacity of adsorbents to adsorb; however, adsorption of MB increases as time increases, and this is due to the fact that the occupation of available vacant active sites on the surface of the adsorbents needs enough time to intercalate inside the zeolites. However, as shown in Figure 5, at lower contact time, the removal efficiency of Z-A was higher than that of Z-B because of its small crystal size indicated in the XRD results; however, as time increases, the removal efficiency of both zeolite types shows a comparable result, as shown in Figure 4. At the cost of using expensive autoclave and high temperature that are related to safety issues, the difference in the removal efficiency in the two zeolites was insignificant.Figure 4
Effect of contact time on removal efficiency of Z-A and Z-B.Figure 5
(a) Effect of dye concentration on removal efficiency and (b) adsorption capacity of Z-A and Z-B.
(a)(b)


#### 3.4.2. Effect of Initial Concentrations

Figures5(a) and 5(b) show absorbance of Z-A and Z-B with function wavelength of 200–800 nm and concentration of methylene blue of 10–40 pmm). In both figures, the absorbance at specific wavelength increases as the MB concentration increases, and this is due to large absorption of visible light by MB in the solution. Figure 6 indicates effect of dye concentration on removal efficiency of Z-A and Z-B. For both zeolites, the removal efficiency of methylene blue decreases as the concentration of methylene blue increases from 10 ppm to 40 ppm because, at low initial concentration of methylene blue, there are many adsorption sites, so that there is high removal efficiency, but as the dyes’ concentration increases, all the active sites of both zeolites are occupied; thus there is little removal efficiency. Figure 6 shows the impact of the initial concentration of the dyes on the adsorption capacity (qe (mg/g)). As indicated in this figure, the amount of absorbed dyes (mg) per g increases as the initial concentration of the dyes increases because the active sites of Z-A and Z-B adsorbents were enclosed by much more methylene blue. Thus, adsorbed MB is increased due to the formation of gradient concentration between aqueous solution and adsorbent surface [37, 39, 40]. Here also the MB removal efficiencies of Z-A and Z-B were comparable.Figure 6
Effect of adsorbent dosage on removal efficiency of Z-A and Z-B.Table2 shows relative absorption capacities of various absorbents. As indicated in Table 1, the absorption capacities of zeolite Linde Type A and sodalite octahydrate zeolite were comparable.Table 2
Relative absorption capacities of various absorbents.
DyeAbsorbentAbsorption capacity (mg/g)ReferenceMBBagasse ash4.41[41]MBNatural zeolite1.5[42]MBZeolite Linde Type A3.9This workMBSodalite octahydrate zeolite3.5This workMBZSM-5 zeolite4.31[43]MBRice husk6[44]MBActivated carbon4.59[45]


#### 3.4.3. Effect of Mass of Adsorbent

Figure6 shows the absorbance of Z-A and Z-B with function of wavelength and mass of absorbent (1–5 g). The absorbances of both zeolites showed the same trends; that is, as mass of absorbent increases, the absorbance decreases because more MB is removed by the zeolites. Figure 7 indicates the removal efficiency of Z-A and Z-B with function mass absorbents. The removal efficiency of both zeolites increases with the increases of its mass because of more adsorption sites introduced to the solution. By increasing the number of adsorbent particles, there is more methylene blue attached, also reported in [46, 47]. The removal efficiencies of both zeolites were also comparable.Figure 7
Effect of pH on removal efficiency of Z-A and Z-B.


#### 3.4.4. Effect of pH Value

Figure7 shows the effect of pH of the solution on the removal efficiency Z-A and Z-B. The results revealed that the removal efficiencies of both zeolites were significantly dependent on the pH values because at low pH value the adsorbent shows less removal efficiency due to the fact that highly acidic hydrogen ion competes with MB ions to occupy the sites of the adsorbent [48], and at higher pH value, which is a high alkali region, there is a generation of more hydroxyl sites, which also decreases the adsorption efficiency of the material [49]. So the optimum pH value was obtained to be 7. Here also the removal efficiencies of both zeolites have minor differences.


#### 3.4.5. Recyclability Study

Figures8(a) and 8(b) show recyclability of Z-A and Z-B for three time tests. In both figures, the blue bar indicates that after absorbent recovered from solution by filtering, it was washed by distilled water, and red bar indicates the recovered absorbent washed by strong acid. The results showed the amount of MB adsorbed onto Z-A and Z-B for the three cycles, by using different regeneration techniques [50, 51]. The first method was done by washing the adsorbents using distilled water several times; then they were centrifuged and finally oven-dried. In the cause of using this method, it was difficult to regenerate the samples because of the MB ion triggered inside the pores of adsorbent. The other method of reusability done on this work was using of strong acid (0.005 mol of HCl) to wash the adsorbent, after which it was centrifuged and dried. The result indicated that the prepared adsorbent can be easily regenerated by using acid rather than water, so acid treated samples have better reusability and stability than those which were washed by water. In both cases, Z-A and Z-B show the same trends of recyclability.Figure 8
(a) The recyclability of Z-A and (b) the recyclability of Z-B.
(a)(b)


### 3.5. Adsorption Isotherm Analysis

The equilibrium model used to characterize equilibrium behavior by describing the amount of adsorbate adsorbed as a function of gas or liquid at a constant temperature is called an isotherm [52]. The basis of the study of the adsorption process is the adsorption isotherm model. The adsorption isotherm is the main source for measuring the adsorption of substances. Several equilibrium models have been used to explain the adsorption process, such as Langmuir, Freundlich, Temkin, Slips, Hill, Radke-Prausnitz, and Flory-Huggins isotherms. Langmuir and Freundlich adsorption isotherms are the most commonly used models [52]. The Langmuir model assumes that the surface of the adsorbent for removing adsorbate is uniform and flat, and the adsorbed molecules or ions have no interaction. On the other hand, the Freundlich model assumes that adsorption occurs on an uneven surface, each local adsorption site has its own binding energy, and the strongest binding site is established before the end of the adsorption process [53, 54]. Among them, qe (mg/g) is the equilibrium adsorption capacity, qmax (mg/g) is the maximum adsorption capacity that the adsorbent can achieve, l/mg is the Langmuir adsorption constant, Ce (mg/l) is the equilibrium concentration, KF is Freund's constant ((mg/g) (l/g)1/n), which represents the binding energy of the adsorbent, and n is the adsorption intensity.


#### 3.5.1. Langmuir Isotherm Model

Langmuir isotherm on solid surface assumes that adsorption is homogeneous or monolayer adsorption [55]. The Langmuir isotherm model can be described according to the equation in Table 3 to determine adsorption isotherm.Table 3
Langmuir and Freundlich parameters extracted from isotherm models.
Isotherm modelNonlinear equationLinear equationPlotReferenceLangmuirqe = qmaxbc/(1 + bCe)Ce/qe=1/bqmax+Ce/qeCe/qevsCeSwensnet et al., [56]Freundlichqe = KFCe1/nlnqe=1/nlnCe+lnkFlnqevslnCe[36]


#### 3.5.2. Freundlich Isotherm Model

Freundlich isotherm is an empirical relation between the concentrations of a solute on the surface of an adsorbent to the concentration of the solute present in the liquid [35]. The Freundlich isotherm equation is described in Table 1. Therefore, based on the results shown in Figures 9 and 10, the adsorption data conformed to both the Langmuir and Freundlich isotherms. Both isotherms were apparently best fitted models with a correlation coefficient (R2) greater than 98; thus, this proves the adsorption of methylene blue on surfaces of Z-A and Z-B due to porous nature of the adsorbent [35]. The maximum adsorption capacity of Z-A towards methylene blue was 4.0299 mg/g and it was 3.678 mg/g for Z-B samples as determined using Langmuir equation. Furthermore, the dimensionless Freundlich constant n > 1, which implies adsorption was favorable on heterogeneous surface. Generally, the value of 1/n < 0.5 implies easy adsorption, and 1/n > 2 implies difficult adsorption [36]. In this work, the adsorption intensity (1/n) = 0.609 and 0.70914 for Z-A and Z-B samples, respectively, and this implied that the adsorption of methylene blue on the surface of zeolite was favorable as the isotherm [57]. So, we can conclude that adsorption was favorable according to both isotherm models.Figure 9
(a) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-A. (b) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-B.
(a)(b)Figure 10
(a) Plot of ln(qe) against ln(Ce) which shows Freundlich isotherm of Z-A. (b) Plot of ln(qe) against ln(Ce) of Freundlich isotherm Z-B.
(a)(b)


### 3.6. Adsorption Kinetics

The study of adsorption kinetics is very important for wastewater purification because it provides important information about process characteristics, reaction forms, and adsorption reaction mechanisms. The adsorption process is mainly described by pseudo-first-order and pseudo-second-order kinetic models based on chemical reaction kinetics [58].


#### 3.6.1. Pseudo-First-Order

Lagergren's pseudo-first-order model is based on the assumption that the rate of change of solute absorption over time is proportional to the difference in saturation concentration and the amount of solid absorbed over time, which is usually applied in the early stages of kinetics. When adsorption occurs through interface diffusion, this pseudo-first-order large particle rate equation is usually observed [57, 58]. The pseudo-first-order Lagergren model can be described by the equation shown in Table 4.Table 4
Equations for pseudo-first-order and pseudo-second-order kinetic models.
Isotherm modelEquationPlotReferencePseudo-first-orderLog (qe − qt) = logqe − K1/2.303 tLog (qe − qt) vs time[59]Pseudo-second-ordert/qt = 1/K2qe2+t/qet/qt vs time[60]According to the pseudo-first-order dynamics, if the adsorption process follows the true first-order kinetics, the intersection of log (qe -qt) and t is equal to the logarithm of qe, which is determined by experiment or corrected to match the experimental value. In this case, the pseudo-first-order Lagergren equation does not fit the entire adsorption time range as shown in Figure 11.Figure 11
Pseudo-first-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


#### 3.6.2. Pseudo-Second-Order

The pseudo-second-order kinetic model is based on the assumption that the rate-limiting step is chemical adsorption or chemisorption and predicts the behavior over the entire adsorption range. In this state, the adsorption rate depends on the adsorption capacity, not the concentration. An important advantage of this model over first-order storage quantities is that the equilibrium adsorption capacity can be calculated from the model; therefore, there is theoretically no need to evaluate the equilibrium adsorption capacity from the experiment [57, 59–62]. In pseudo-second-order kinetics, the adsorption process follows true second-order kinetics as shown in Figure 12; then the intercept of t/qe versus t plots had been adjusted to fit the experimental value, so pseudo-second-order kinetics were fitted well for the whole range of adsorption time. In pseudo-first-order kinetic models, there is a big difference in qe, exp and qe, cal, which indicates that the adsorption of methylene blue by Z-A and Z-B adsorbent does not obey pseudo-first-order kinetics; on the other hand, the values of qe, exp, and qe, cal obtained in pseudo-second-order model are very close. In addition to this, the correlation coefficient (R2) = 0.76105 and 0.37826, respectively, for Z-A and Z-B in pseudo-first-order kinetics, wile the pseudo-second-order kinetics model showed a higher correlation coefficient (R2) of 0.99982 and 0.99522 for Z-A and Z-B, respectively, so we can confirm that the pseudo-second-order model is the best fitted model for the adsorbent.Figure 12
Pseudo-second-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


## 3.1. XRD Analysis

Figure1 shows the XRD results of Z-A and Z-B. In Z-A, the diffraction peaks at angles of 2θ = 13.99°, 19.8°, 24.3°, 31.5°, 34.6°, 37.4°, 42.7°, and 58.09° correspond to the (100), (200), (211), (310), (222), (321), (411), and (440) crystallographic planes, indicating sodalite octahydrate zeolite. These results are aligned with the results reported in [33]. In addition to this, the crystal structure of Z-A was compared with the standard zeolite material and it was found that they have similar structures. The XRD pattern of Z-B was also indicated in Figure 1. Diffraction peaks of Z-B were observed at angles of 2θ = 7.3°, 10.3°, 12.6°, 16.2°, 21.7°, 24.1°, 27.2°, 30°, and 34.2° corresponding to the (100), (001) (101), (110), (101), (101), (210), (001), and (220) crystallographic planes. These data revealed that Z-B is zeolite LTA that is matched with standard of zeolite LTA. We can conclude that the synthesis of zeolite by using different techniques may give different types of zeolite. But, in both cases of zeolites, no considerable amounts of amorphous phases were observed. Here, even though the type of nonhydrothermally synthesized zeolite has different crystal structure as compared with hydrothermal synthesized zeolite, obtaining zeolite without using expensive autoclave is important in order to reduce the cost and for safety issue.Figure 1
XRD patterns of Z-A and Z-B.The crystalline sizes of Z-A and Z-B particles were calculated to be around 23.14 nm and 49.73 nm, respectively, using the Scherrer formula, according to the following equation [34]:(3)D=Kλβcosθ,where λ = 0.15406 nm (wavelength of X-ray), K = 0.9 (Scherrer constant), D is crystalline size (nm), and = FWHM (full width at half maximum). As is expected, the crystal size of hydrothermally synthesized zeolite was smaller than that of nonhydrothermally synthesized one because of the high pressure in the autoclave. Practically, as particle size decreases, the surface area of the material increases, which promotes the absorption of the pollutants in the wastewater [31]. The effect of crystal size on the removal efficiency will be addressed in Section 3.2.


## 3.2. SEM Analysis

Figure2 shows the morphologies of the synthesized zeolites Z-A and Z-B at different magnification. Z-A shows interconnected porous structure with an average particle size less than 24 nm as indicated in Figure 2(a). Similarly, porous structure having interconnection was observed in case Z-B as shown in Figure 2(b). In both cases, spherical-like structures were observed at magnification of 1000x. For the distinct zeolite types, it is reported that different crystal shapes of zeolite are probably associated with different raw materials and preparation methods applied ([32], [33]). Here, although the synthesizing methods were different, the precursors were the same for both zeolites. As a result, a cluster of spherical shapes were observed in both zeolites.Figure 2
SEM image of (a) Z-A and (b) Z-B at different magnification.
(a)(b)


## 3.3. FTIR Analysis

The FTIR spectra of Z-A and and Z-B are shown in Figure3. The FTIR results showed almost similar spectra of packs with different intensity. In the fingerprint region, the spectrum shows a broad and intense band around 1000 cm−1, showing that the characteristic of antisymmetric stretching vibration of the Si-O-Si and a less intense band around 789 cm−1 is due to Si-O-Si symmetric stretching vibrations; this has the same stretching as reported in [35]. Both synthesized zeolites show peaks around 3500 cm−1 and 1650 cm−1, which corresponds to the stretching of the hydroxyl groups (OH-) from adsorbed water molecules. The sharpness of the peaks could determine the amount of water (H2O) molecule adsorbed on the surface of the samples, which in turn can be used to estimate the amount of the hydroxyl functional groups present in the sample as indicated in other research [34]. Other functional groups that exist in the samples are Si-O-Si, indicated by the absorption band at 1020 cm−1, Al-O indicated by peak at 770 cm−1, and Si-O-Al, indicated by the presence of an absorption band at the region between 420 and 494 cm−1. Concerning the existence of functional groups Si, O, and Al observed in the spectrum, it was then concluded that the formation of the zeolite framework was confirmed by the FTIR results as indicated in [36].Figure 3
FTIR spectra of Z-A and Z-B.


## 3.4. Adsorption Performance Analysis

The efficiency of adsorbent in water purification is affected by many factors such as contact time, initial concentration of impurity, adsorbent dose, and pH [37, 38]. In this section, the adsorbent performance is studied with respect to those parameters. Similar to the intrinsic properties of the absorbents, the external factors mentioned have a significant impact on the absorbing properties of absorbents.


### 3.4.1. Effect of Contact Time on Removal Efficiency of Absorbent

Figure4 shows the effect of the contact time on the MB removal efficiency by both Z-A and Z-B. The adsorption of methylene blue by Z-A and Z-B at room temperature, initial methylene blue concentration of 10 ppm, and adsorbent dosage (Z-A and Z-B) of 3 g/L were used to study the effect of contact time for 150 minutes on the removal efficiency of adsorbents. The adsorptions of methylene blue were slow for both Z-A and Z-B with respect to time because of the low capacity of adsorbents to adsorb; however, adsorption of MB increases as time increases, and this is due to the fact that the occupation of available vacant active sites on the surface of the adsorbents needs enough time to intercalate inside the zeolites. However, as shown in Figure 5, at lower contact time, the removal efficiency of Z-A was higher than that of Z-B because of its small crystal size indicated in the XRD results; however, as time increases, the removal efficiency of both zeolite types shows a comparable result, as shown in Figure 4. At the cost of using expensive autoclave and high temperature that are related to safety issues, the difference in the removal efficiency in the two zeolites was insignificant.Figure 4
Effect of contact time on removal efficiency of Z-A and Z-B.Figure 5
(a) Effect of dye concentration on removal efficiency and (b) adsorption capacity of Z-A and Z-B.
(a)(b)


### 3.4.2. Effect of Initial Concentrations

Figures5(a) and 5(b) show absorbance of Z-A and Z-B with function wavelength of 200–800 nm and concentration of methylene blue of 10–40 pmm). In both figures, the absorbance at specific wavelength increases as the MB concentration increases, and this is due to large absorption of visible light by MB in the solution. Figure 6 indicates effect of dye concentration on removal efficiency of Z-A and Z-B. For both zeolites, the removal efficiency of methylene blue decreases as the concentration of methylene blue increases from 10 ppm to 40 ppm because, at low initial concentration of methylene blue, there are many adsorption sites, so that there is high removal efficiency, but as the dyes’ concentration increases, all the active sites of both zeolites are occupied; thus there is little removal efficiency. Figure 6 shows the impact of the initial concentration of the dyes on the adsorption capacity (qe (mg/g)). As indicated in this figure, the amount of absorbed dyes (mg) per g increases as the initial concentration of the dyes increases because the active sites of Z-A and Z-B adsorbents were enclosed by much more methylene blue. Thus, adsorbed MB is increased due to the formation of gradient concentration between aqueous solution and adsorbent surface [37, 39, 40]. Here also the MB removal efficiencies of Z-A and Z-B were comparable.Figure 6
Effect of adsorbent dosage on removal efficiency of Z-A and Z-B.Table2 shows relative absorption capacities of various absorbents. As indicated in Table 1, the absorption capacities of zeolite Linde Type A and sodalite octahydrate zeolite were comparable.Table 2
Relative absorption capacities of various absorbents.
DyeAbsorbentAbsorption capacity (mg/g)ReferenceMBBagasse ash4.41[41]MBNatural zeolite1.5[42]MBZeolite Linde Type A3.9This workMBSodalite octahydrate zeolite3.5This workMBZSM-5 zeolite4.31[43]MBRice husk6[44]MBActivated carbon4.59[45]


### 3.4.3. Effect of Mass of Adsorbent

Figure6 shows the absorbance of Z-A and Z-B with function of wavelength and mass of absorbent (1–5 g). The absorbances of both zeolites showed the same trends; that is, as mass of absorbent increases, the absorbance decreases because more MB is removed by the zeolites. Figure 7 indicates the removal efficiency of Z-A and Z-B with function mass absorbents. The removal efficiency of both zeolites increases with the increases of its mass because of more adsorption sites introduced to the solution. By increasing the number of adsorbent particles, there is more methylene blue attached, also reported in [46, 47]. The removal efficiencies of both zeolites were also comparable.Figure 7
Effect of pH on removal efficiency of Z-A and Z-B.


### 3.4.4. Effect of pH Value

Figure7 shows the effect of pH of the solution on the removal efficiency Z-A and Z-B. The results revealed that the removal efficiencies of both zeolites were significantly dependent on the pH values because at low pH value the adsorbent shows less removal efficiency due to the fact that highly acidic hydrogen ion competes with MB ions to occupy the sites of the adsorbent [48], and at higher pH value, which is a high alkali region, there is a generation of more hydroxyl sites, which also decreases the adsorption efficiency of the material [49]. So the optimum pH value was obtained to be 7. Here also the removal efficiencies of both zeolites have minor differences.


### 3.4.5. Recyclability Study

Figures8(a) and 8(b) show recyclability of Z-A and Z-B for three time tests. In both figures, the blue bar indicates that after absorbent recovered from solution by filtering, it was washed by distilled water, and red bar indicates the recovered absorbent washed by strong acid. The results showed the amount of MB adsorbed onto Z-A and Z-B for the three cycles, by using different regeneration techniques [50, 51]. The first method was done by washing the adsorbents using distilled water several times; then they were centrifuged and finally oven-dried. In the cause of using this method, it was difficult to regenerate the samples because of the MB ion triggered inside the pores of adsorbent. The other method of reusability done on this work was using of strong acid (0.005 mol of HCl) to wash the adsorbent, after which it was centrifuged and dried. The result indicated that the prepared adsorbent can be easily regenerated by using acid rather than water, so acid treated samples have better reusability and stability than those which were washed by water. In both cases, Z-A and Z-B show the same trends of recyclability.Figure 8
(a) The recyclability of Z-A and (b) the recyclability of Z-B.
(a)(b)


## 3.4.1. Effect of Contact Time on Removal Efficiency of Absorbent

Figure4 shows the effect of the contact time on the MB removal efficiency by both Z-A and Z-B. The adsorption of methylene blue by Z-A and Z-B at room temperature, initial methylene blue concentration of 10 ppm, and adsorbent dosage (Z-A and Z-B) of 3 g/L were used to study the effect of contact time for 150 minutes on the removal efficiency of adsorbents. The adsorptions of methylene blue were slow for both Z-A and Z-B with respect to time because of the low capacity of adsorbents to adsorb; however, adsorption of MB increases as time increases, and this is due to the fact that the occupation of available vacant active sites on the surface of the adsorbents needs enough time to intercalate inside the zeolites. However, as shown in Figure 5, at lower contact time, the removal efficiency of Z-A was higher than that of Z-B because of its small crystal size indicated in the XRD results; however, as time increases, the removal efficiency of both zeolite types shows a comparable result, as shown in Figure 4. At the cost of using expensive autoclave and high temperature that are related to safety issues, the difference in the removal efficiency in the two zeolites was insignificant.Figure 4
Effect of contact time on removal efficiency of Z-A and Z-B.Figure 5
(a) Effect of dye concentration on removal efficiency and (b) adsorption capacity of Z-A and Z-B.
(a)(b)


## 3.4.2. Effect of Initial Concentrations

Figures5(a) and 5(b) show absorbance of Z-A and Z-B with function wavelength of 200–800 nm and concentration of methylene blue of 10–40 pmm). In both figures, the absorbance at specific wavelength increases as the MB concentration increases, and this is due to large absorption of visible light by MB in the solution. Figure 6 indicates effect of dye concentration on removal efficiency of Z-A and Z-B. For both zeolites, the removal efficiency of methylene blue decreases as the concentration of methylene blue increases from 10 ppm to 40 ppm because, at low initial concentration of methylene blue, there are many adsorption sites, so that there is high removal efficiency, but as the dyes’ concentration increases, all the active sites of both zeolites are occupied; thus there is little removal efficiency. Figure 6 shows the impact of the initial concentration of the dyes on the adsorption capacity (qe (mg/g)). As indicated in this figure, the amount of absorbed dyes (mg) per g increases as the initial concentration of the dyes increases because the active sites of Z-A and Z-B adsorbents were enclosed by much more methylene blue. Thus, adsorbed MB is increased due to the formation of gradient concentration between aqueous solution and adsorbent surface [37, 39, 40]. Here also the MB removal efficiencies of Z-A and Z-B were comparable.Figure 6
Effect of adsorbent dosage on removal efficiency of Z-A and Z-B.Table2 shows relative absorption capacities of various absorbents. As indicated in Table 1, the absorption capacities of zeolite Linde Type A and sodalite octahydrate zeolite were comparable.Table 2
Relative absorption capacities of various absorbents.
DyeAbsorbentAbsorption capacity (mg/g)ReferenceMBBagasse ash4.41[41]MBNatural zeolite1.5[42]MBZeolite Linde Type A3.9This workMBSodalite octahydrate zeolite3.5This workMBZSM-5 zeolite4.31[43]MBRice husk6[44]MBActivated carbon4.59[45]


## 3.4.3. Effect of Mass of Adsorbent

Figure6 shows the absorbance of Z-A and Z-B with function of wavelength and mass of absorbent (1–5 g). The absorbances of both zeolites showed the same trends; that is, as mass of absorbent increases, the absorbance decreases because more MB is removed by the zeolites. Figure 7 indicates the removal efficiency of Z-A and Z-B with function mass absorbents. The removal efficiency of both zeolites increases with the increases of its mass because of more adsorption sites introduced to the solution. By increasing the number of adsorbent particles, there is more methylene blue attached, also reported in [46, 47]. The removal efficiencies of both zeolites were also comparable.Figure 7
Effect of pH on removal efficiency of Z-A and Z-B.


## 3.4.4. Effect of pH Value

Figure7 shows the effect of pH of the solution on the removal efficiency Z-A and Z-B. The results revealed that the removal efficiencies of both zeolites were significantly dependent on the pH values because at low pH value the adsorbent shows less removal efficiency due to the fact that highly acidic hydrogen ion competes with MB ions to occupy the sites of the adsorbent [48], and at higher pH value, which is a high alkali region, there is a generation of more hydroxyl sites, which also decreases the adsorption efficiency of the material [49]. So the optimum pH value was obtained to be 7. Here also the removal efficiencies of both zeolites have minor differences.


## 3.4.5. Recyclability Study

Figures8(a) and 8(b) show recyclability of Z-A and Z-B for three time tests. In both figures, the blue bar indicates that after absorbent recovered from solution by filtering, it was washed by distilled water, and red bar indicates the recovered absorbent washed by strong acid. The results showed the amount of MB adsorbed onto Z-A and Z-B for the three cycles, by using different regeneration techniques [50, 51]. The first method was done by washing the adsorbents using distilled water several times; then they were centrifuged and finally oven-dried. In the cause of using this method, it was difficult to regenerate the samples because of the MB ion triggered inside the pores of adsorbent. The other method of reusability done on this work was using of strong acid (0.005 mol of HCl) to wash the adsorbent, after which it was centrifuged and dried. The result indicated that the prepared adsorbent can be easily regenerated by using acid rather than water, so acid treated samples have better reusability and stability than those which were washed by water. In both cases, Z-A and Z-B show the same trends of recyclability.Figure 8
(a) The recyclability of Z-A and (b) the recyclability of Z-B.
(a)(b)


## 3.5. Adsorption Isotherm Analysis

The equilibrium model used to characterize equilibrium behavior by describing the amount of adsorbate adsorbed as a function of gas or liquid at a constant temperature is called an isotherm [52]. The basis of the study of the adsorption process is the adsorption isotherm model. The adsorption isotherm is the main source for measuring the adsorption of substances. Several equilibrium models have been used to explain the adsorption process, such as Langmuir, Freundlich, Temkin, Slips, Hill, Radke-Prausnitz, and Flory-Huggins isotherms. Langmuir and Freundlich adsorption isotherms are the most commonly used models [52]. The Langmuir model assumes that the surface of the adsorbent for removing adsorbate is uniform and flat, and the adsorbed molecules or ions have no interaction. On the other hand, the Freundlich model assumes that adsorption occurs on an uneven surface, each local adsorption site has its own binding energy, and the strongest binding site is established before the end of the adsorption process [53, 54]. Among them, qe (mg/g) is the equilibrium adsorption capacity, qmax (mg/g) is the maximum adsorption capacity that the adsorbent can achieve, l/mg is the Langmuir adsorption constant, Ce (mg/l) is the equilibrium concentration, KF is Freund's constant ((mg/g) (l/g)1/n), which represents the binding energy of the adsorbent, and n is the adsorption intensity.


### 3.5.1. Langmuir Isotherm Model

Langmuir isotherm on solid surface assumes that adsorption is homogeneous or monolayer adsorption [55]. The Langmuir isotherm model can be described according to the equation in Table 3 to determine adsorption isotherm.Table 3
Langmuir and Freundlich parameters extracted from isotherm models.
Isotherm modelNonlinear equationLinear equationPlotReferenceLangmuirqe = qmaxbc/(1 + bCe)Ce/qe=1/bqmax+Ce/qeCe/qevsCeSwensnet et al., [56]Freundlichqe = KFCe1/nlnqe=1/nlnCe+lnkFlnqevslnCe[36]


### 3.5.2. Freundlich Isotherm Model

Freundlich isotherm is an empirical relation between the concentrations of a solute on the surface of an adsorbent to the concentration of the solute present in the liquid [35]. The Freundlich isotherm equation is described in Table 1. Therefore, based on the results shown in Figures 9 and 10, the adsorption data conformed to both the Langmuir and Freundlich isotherms. Both isotherms were apparently best fitted models with a correlation coefficient (R2) greater than 98; thus, this proves the adsorption of methylene blue on surfaces of Z-A and Z-B due to porous nature of the adsorbent [35]. The maximum adsorption capacity of Z-A towards methylene blue was 4.0299 mg/g and it was 3.678 mg/g for Z-B samples as determined using Langmuir equation. Furthermore, the dimensionless Freundlich constant n > 1, which implies adsorption was favorable on heterogeneous surface. Generally, the value of 1/n < 0.5 implies easy adsorption, and 1/n > 2 implies difficult adsorption [36]. In this work, the adsorption intensity (1/n) = 0.609 and 0.70914 for Z-A and Z-B samples, respectively, and this implied that the adsorption of methylene blue on the surface of zeolite was favorable as the isotherm [57]. So, we can conclude that adsorption was favorable according to both isotherm models.Figure 9
(a) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-A. (b) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-B.
(a)(b)Figure 10
(a) Plot of ln(qe) against ln(Ce) which shows Freundlich isotherm of Z-A. (b) Plot of ln(qe) against ln(Ce) of Freundlich isotherm Z-B.
(a)(b)


## 3.5.1. Langmuir Isotherm Model

Langmuir isotherm on solid surface assumes that adsorption is homogeneous or monolayer adsorption [55]. The Langmuir isotherm model can be described according to the equation in Table 3 to determine adsorption isotherm.Table 3
Langmuir and Freundlich parameters extracted from isotherm models.
Isotherm modelNonlinear equationLinear equationPlotReferenceLangmuirqe = qmaxbc/(1 + bCe)Ce/qe=1/bqmax+Ce/qeCe/qevsCeSwensnet et al., [56]Freundlichqe = KFCe1/nlnqe=1/nlnCe+lnkFlnqevslnCe[36]


## 3.5.2. Freundlich Isotherm Model

Freundlich isotherm is an empirical relation between the concentrations of a solute on the surface of an adsorbent to the concentration of the solute present in the liquid [35]. The Freundlich isotherm equation is described in Table 1. Therefore, based on the results shown in Figures 9 and 10, the adsorption data conformed to both the Langmuir and Freundlich isotherms. Both isotherms were apparently best fitted models with a correlation coefficient (R2) greater than 98; thus, this proves the adsorption of methylene blue on surfaces of Z-A and Z-B due to porous nature of the adsorbent [35]. The maximum adsorption capacity of Z-A towards methylene blue was 4.0299 mg/g and it was 3.678 mg/g for Z-B samples as determined using Langmuir equation. Furthermore, the dimensionless Freundlich constant n > 1, which implies adsorption was favorable on heterogeneous surface. Generally, the value of 1/n < 0.5 implies easy adsorption, and 1/n > 2 implies difficult adsorption [36]. In this work, the adsorption intensity (1/n) = 0.609 and 0.70914 for Z-A and Z-B samples, respectively, and this implied that the adsorption of methylene blue on the surface of zeolite was favorable as the isotherm [57]. So, we can conclude that adsorption was favorable according to both isotherm models.Figure 9
(a) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-A. (b) Langmuir isotherm: a plot of 1/qe against 1/Ce of sample Z-B.
(a)(b)Figure 10
(a) Plot of ln(qe) against ln(Ce) which shows Freundlich isotherm of Z-A. (b) Plot of ln(qe) against ln(Ce) of Freundlich isotherm Z-B.
(a)(b)


## 3.6. Adsorption Kinetics

The study of adsorption kinetics is very important for wastewater purification because it provides important information about process characteristics, reaction forms, and adsorption reaction mechanisms. The adsorption process is mainly described by pseudo-first-order and pseudo-second-order kinetic models based on chemical reaction kinetics [58].


### 3.6.1. Pseudo-First-Order

Lagergren's pseudo-first-order model is based on the assumption that the rate of change of solute absorption over time is proportional to the difference in saturation concentration and the amount of solid absorbed over time, which is usually applied in the early stages of kinetics. When adsorption occurs through interface diffusion, this pseudo-first-order large particle rate equation is usually observed [57, 58]. The pseudo-first-order Lagergren model can be described by the equation shown in Table 4.Table 4
Equations for pseudo-first-order and pseudo-second-order kinetic models.
Isotherm modelEquationPlotReferencePseudo-first-orderLog (qe − qt) = logqe − K1/2.303 tLog (qe − qt) vs time[59]Pseudo-second-ordert/qt = 1/K2qe2+t/qet/qt vs time[60]According to the pseudo-first-order dynamics, if the adsorption process follows the true first-order kinetics, the intersection of log (qe -qt) and t is equal to the logarithm of qe, which is determined by experiment or corrected to match the experimental value. In this case, the pseudo-first-order Lagergren equation does not fit the entire adsorption time range as shown in Figure 11.Figure 11
Pseudo-first-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


### 3.6.2. Pseudo-Second-Order

The pseudo-second-order kinetic model is based on the assumption that the rate-limiting step is chemical adsorption or chemisorption and predicts the behavior over the entire adsorption range. In this state, the adsorption rate depends on the adsorption capacity, not the concentration. An important advantage of this model over first-order storage quantities is that the equilibrium adsorption capacity can be calculated from the model; therefore, there is theoretically no need to evaluate the equilibrium adsorption capacity from the experiment [57, 59–62]. In pseudo-second-order kinetics, the adsorption process follows true second-order kinetics as shown in Figure 12; then the intercept of t/qe versus t plots had been adjusted to fit the experimental value, so pseudo-second-order kinetics were fitted well for the whole range of adsorption time. In pseudo-first-order kinetic models, there is a big difference in qe, exp and qe, cal, which indicates that the adsorption of methylene blue by Z-A and Z-B adsorbent does not obey pseudo-first-order kinetics; on the other hand, the values of qe, exp, and qe, cal obtained in pseudo-second-order model are very close. In addition to this, the correlation coefficient (R2) = 0.76105 and 0.37826, respectively, for Z-A and Z-B in pseudo-first-order kinetics, wile the pseudo-second-order kinetics model showed a higher correlation coefficient (R2) of 0.99982 and 0.99522 for Z-A and Z-B, respectively, so we can confirm that the pseudo-second-order model is the best fitted model for the adsorbent.Figure 12
Pseudo-second-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


## 3.6.1. Pseudo-First-Order

Lagergren's pseudo-first-order model is based on the assumption that the rate of change of solute absorption over time is proportional to the difference in saturation concentration and the amount of solid absorbed over time, which is usually applied in the early stages of kinetics. When adsorption occurs through interface diffusion, this pseudo-first-order large particle rate equation is usually observed [57, 58]. The pseudo-first-order Lagergren model can be described by the equation shown in Table 4.Table 4
Equations for pseudo-first-order and pseudo-second-order kinetic models.
Isotherm modelEquationPlotReferencePseudo-first-orderLog (qe − qt) = logqe − K1/2.303 tLog (qe − qt) vs time[59]Pseudo-second-ordert/qt = 1/K2qe2+t/qet/qt vs time[60]According to the pseudo-first-order dynamics, if the adsorption process follows the true first-order kinetics, the intersection of log (qe -qt) and t is equal to the logarithm of qe, which is determined by experiment or corrected to match the experimental value. In this case, the pseudo-first-order Lagergren equation does not fit the entire adsorption time range as shown in Figure 11.Figure 11
Pseudo-first-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


## 3.6.2. Pseudo-Second-Order

The pseudo-second-order kinetic model is based on the assumption that the rate-limiting step is chemical adsorption or chemisorption and predicts the behavior over the entire adsorption range. In this state, the adsorption rate depends on the adsorption capacity, not the concentration. An important advantage of this model over first-order storage quantities is that the equilibrium adsorption capacity can be calculated from the model; therefore, there is theoretically no need to evaluate the equilibrium adsorption capacity from the experiment [57, 59–62]. In pseudo-second-order kinetics, the adsorption process follows true second-order kinetics as shown in Figure 12; then the intercept of t/qe versus t plots had been adjusted to fit the experimental value, so pseudo-second-order kinetics were fitted well for the whole range of adsorption time. In pseudo-first-order kinetic models, there is a big difference in qe, exp and qe, cal, which indicates that the adsorption of methylene blue by Z-A and Z-B adsorbent does not obey pseudo-first-order kinetics; on the other hand, the values of qe, exp, and qe, cal obtained in pseudo-second-order model are very close. In addition to this, the correlation coefficient (R2) = 0.76105 and 0.37826, respectively, for Z-A and Z-B in pseudo-first-order kinetics, wile the pseudo-second-order kinetics model showed a higher correlation coefficient (R2) of 0.99982 and 0.99522 for Z-A and Z-B, respectively, so we can confirm that the pseudo-second-order model is the best fitted model for the adsorbent.Figure 12
Pseudo-second-order model for (a) Z-A and (b) Z-B samples.
(a)(b)


## 4. Conclusions

Sodalite octahydrate zeolite is an alternative absorbent for removing methylene blue from aqueous solution with respect to zeolite LTA because sodalite octahydrate zeolite was synthesized by solution method, which is a cheap, easy, and safe method, and zeolite LTA was synthesized by the hydrothermal method that requires expensive autoclave. The absorption capacities of sodalite octahydrate zeolite and zeolite LTA were 3.5 mg/g and 3.9 mg/g at 40 mg/L, respectively. The pH values of the solution affect the absorption capacity of both zeolites. The optimum removal efficiencies of both zeolites were observed at pH value of 7. The removal efficiencies of mentioned zeolites increase with increasing of absorbent dosage because absorption active site increases with the scale-up of the absorbent dosage. The stabilities of both absorbents were checked three times and it was found that they were stable. The absorption isotherms of sodalite octahydrate zeolite and zeolite LTA were effectively fitted with the Freundlich and Langmuir models. Moreover, the absorption kinetics of both zeolites follow the pseudo-second-order kinetics [62]



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