Adsorption of Tetracycline by Magnetic Mesoporous Silica Derived from Bottom Ash—Biomass Power Plant

Abstract

In recent years, the contamination of the aquatic environment with antibiotics, including tetracyclines, has drawn much attention. Bottom ash (BA), a residue from the biomass power plant, was used to synthesize the magnetic mesoporous silica (MMS) and was utilized as an adsorbent for tetracycline (TC) removal from aqueous solutions. The MMS was characterized by Fourier transform-infrared (FTIR), X-ray diffraction (XRD) pattern, and scanning electron microscopy (SEM). Optimum conditions were obtained in overnight incubation at 60 °C, a pH of 6–8, and an adsorption capacity of 276.74 mg/g. The isotherm and kinetic equations pointed to a Langmuir isotherm model and pseudo-first-order kinetic optimum fitting models. Based on the very low values of entropy changes (ΔS°), the negative value of enthalpy changes (ΔH°) (−15.94 kJ/mol), and the negative Gibbs free-energy changes (ΔG°), the adsorption process was physisorption and spontaneous.

1. Introduction

Antibiotics have been commonly used for the prevention and treatment of infectious diseases, both in humans and animals. According to Cycon et al., the situation analyzed in 75 countries, indicates that the antibiotics usage climbed by 65%, and the prediction is that in 2030, the consumption of antibiotics will be higher by 200% compared with 2015 [1]. Among antibiotics, tetracycline (TC) is a conventional, inexpensive antibiotic having wide-ranging antibacterial action that is frequently used in both human and veterinary medicine to prevent infection [2]. Additionally, TC antibiotics are utilized as a growth stimulant for animals and for agricultural purposes [3,4]. Moreover, TC constitutes one of the most important antibiotic families, ranking second in production and usage worldwide [5,6]. Because of its high aqueous solubility, TC has been found in ecological communities such as surface water and groundwater (ranging from 5.4 to 8.1 ng/L) [7], municipal solid waste (greater than 100 ng/L) [8], and soil (ranging from 86.2 to 198.7 µg/kg) [9], and it can be easily transferred to other environments via aqueous matrixes [10], and like other antibiotics, it has a structural framework called naphthol ring that makes it difficult to degrade [11]. Thus, Cycoń et al. [1] suggested that “the reason is that antibiotics are not completely metabolized by humans and animals, and a large proportion of the administered drug is released as the parent compound through feces and urine, discharging into domestic wastewater and into the pits where slurries/sludges are deposited”. Gu and Karthikeyan [12] also found that “in the statewide survey of wastewater treatment plants, the compound TC was the most frequently detected antibiotic (among 25 antibiotics), being present in 80% of the wastewater influent and effluent samples”. However, standard aqueous solution water treatment and spontaneous biodegradation are ineffective in removing TC from aqueous solution [12,13]. Therefore, techniques for the secure and efficient removal of TC from liquid have received a great deal of interest. Methods for TC removal include oxidation [14], photo electrocatalytic [15], degradation [16], membrane processing [17], adsorption [18], permeation [19], flocculation [19], ozonation oxidation [19]. Among these, adsorption has several advantages over other procedures, including ease of operation, inexpensive operating costs, and significant removal efficiency at extremely low TC content in wastewater and water [20]. Particularly, no generation of toxicity of intermediated and by-products during adsorption makes it a more attractive appealing method of treating TC [21]. Until now, some studies investigated the adsorption and removal of tetracyclines on several materials such as graphene oxide [22], activated carbon [23], kaolinite [24], single-walled carbon nanotubes, and multi-walled carbon nanotubes [13]. Carbon nanotubes and graphene, which have a high graphite structure, have a strong TC elimination capacity [25]. However, the high manufacturing, disposal, and regeneration costs of the aforementioned materials would pose significant barriers in their practical application for TC adsorption, as well as the capacity for large-scale application. As a result, there is still a considerable desire for the development of efficient and affordable, simple-to-operate, high-selectivity devices.

Mesoporous silicas are often employed in a variety of applications, including catalysis, separation, drug administration, chemical sensing, optic and electrical devices, rubber reinforcement, and as a template in the production of nanomaterials [26]. Because they possess a large surface area, big pore size, pore volume, and regular channel-type structures, these properties are potential advantages that suit the adsorbate. A renewed interest in the design of adsorbents [27] and catalysts has been sparked by the discovery of hexagonally organized mesoporous silica [28], which has a distinctively large surface area, well-defined pore size, and pore-shaped pores. For the production of surfactant-template silica materials, which need an organic structure-directing agent or template is typically a single surfactant, such as: the Pluronic-type surfactant; poly(ethylene oxide)-b-poly(propylene oxide)-b-poly(ethylene oxide)-Pluronic P123 (EO₂₀PO₇₀EO₂₀) or Pluronic F127 (EO₁₀₆PO₆₀EO₁₀₆) and cetyltrimethylammonium bromide (CTMAB) [29,30,31,32]. Some studies have magnetized the silicates to allow for rapidly and effectively separating the adsorbents from the aqueous solution by an external magnetic field, and incorporating magnetic elements within or on mesoporous silicates has also enhanced the adsorption properties [33].

Some previous researches have investigated magnetic mesoporous silica (MMS) with a variety of structures: including embedded [31], core-shell [32], and yolk-shell [34]. The MMS could be synthesized by combining the mesoporous silica (MS) and magnetic particles. Tetraethyl orthosilicate (TEOS) was used as the silica source, and magnetic particles were prepared from iron (II) chloride tetrahydrate (FeCl₂,4H₂O), iron (III) chloride hexahydrate (FeCl₃·6H₂O) [35]. Because TEOS is expensive and magnetic particles require time and chemicals to make, searches for alternative cheap source are necessary.

Bottom ash (BA)—biomass power plant was a byproduct from the combusting of agricultural waste to produce energy. Approximately 85–95 percent of BA was generated from biomass power plants [36], with SiO₂ making up the majority of this ash [37]. Thus, these residues were an ideal silica source. Currently, Thailand has an abundance of BA—biomass power plants that may be utilized to create magnetic mesoporous silica as an appropriate supply. Generally, utilizing the residues from incineration and combustion plant to produce silica materials is feasible. Some studies have been successful in converting this byproduct to zeolite in alkali solution [38], synthesizing zeolitic material and successfully separating SiO₂ from municipal solid waste incineration (MSWI) ash [39], MCM-41, SBA-15, and SBA-16 mesoporous silica from power plant bottom ash were successfully synthesized for the first time in 2007 [40]. Mesoporous silica was prepared by the sol-gel method from municipal solid waste incineration bottom ash [41] and industrial fly ash [42]. This application would increase the usage of BA—biomass power plants, minimize pollution, and effectively improve the quality of the wastewater process. However, according to our knowledge, there has not been any published research on the synthesis of magnetic silica nanoparticles derived from BA for specific applications in the adsorption of tetracycline.

Therefore, the purpose of this study is to apply magnetic mesoporous silica synthesized from biomass power plant ash to absorb TC from aqueous solutions. The properties of the MMS including surface area, function groups, and crystal structure were analyzed. The research was conducted on the isotherm model, kinetics model, and thermodynamic characteristics of TC’s adsorption onto the MMS. According to the findings of adsorption tests, the MMS demonstrated promising potential for TC removal from water. This study provides a detailed technique for achieving the goal of treating trash with waste in addition to describing how solid waste is utilized as a resource.

2. Materials and Methods

2.1. Materials

In this study, the bottom ash (BA) was obtained from the rubber biomass power plant of Gulf Yala company, Yala, located in southern Thailand. The composition of BA was determined by X-ray fluorescence spectrometer (XRF) as listed in Table S1 (in the Supplementary material). Tetracycline hydrochloride, 96%, Alfa Aesar were used without further purification. Hexadecyltrimethyl ammonium bromide (CTAB) ≥ 98%, Sigma, Cibolo, TX, USA; Sodium hydroxide (NaOH) 98%, Loba Chemie PVT.LTD, India; Hydrochloric acid (HCl) 37%, Qrec, Newzealand; Ethanol 99.9%, Qrec, Newzealand.

2.2. Methods

2.2.1. Synthesis of Magnetic Mesoporous Silica

The biomass power plant ash (BA) which includes 55% SiO₂ content as the main component (Table S1 in the Supplementary Material) was used as raw material for extraction of SiO₂ compound. Briefly, BA sample was ground into small particles (<45 μm) and then washed with DI water several times and dried in the oven at 90 °C for 24 h. The permanent magnets were employed to separate the magnetic and nonmagnetic ashes. The magnetic ash was served as the magnetic component in the magnetic mesoporous silica, and the nonmagnetic ash was utilized to extract silica. The weight ratio between the magnetic and nonmagnetic ashes was 1:10.

A total of 4 g of nonmagnetic ash was refluxed with 100 mL NaOH solution with the concentration 4 M at 90 °C for 16 h in a round bottom flask. Then, the reaction solution was filtered through a membrane of 0.45 µm and the supernatant was collected for adjusting pH at 7 by 5 M HCl and aging at the room temperature for 24 h. The precipitated product was washed with distilled water several times and dried at 90 °C for 24 h. To prepare the sodium silicate solution, the obtained white power and NaOH (4:5 w/w) were precisely weighed and then dissolved in 250 mL of distilled water at 80 °C for 1 day [43]. The white power is SiO₂ and was extracted from BA, which was analyzed by X-ray fluorescence spectrometer (XRF), as shown in Table S2 (in the Supplementary material).

Deionized water (20 mL) was added to 1.2 g hexadecyltrimethyl ammonium bromide (CTAB) and then mixed at 40 °C until a clear solution was observed. The aqueous CTAB solution was slowly mixed with sodium silicate solution and continuously stirred for 15 min. This mixture was added slowly into the magnetic ash; the ratio of the mass of maghemite to Na₂SiO₃ 10% was 1:20, and it was stirred while being heated to 80 °C. After further stirring for 30 min, the pH of the mixture was adjusted to 11 by 5 mol/L HCl solution and then continuously stirred for 6 h. The mixture was kept in a water bath at 80 °C for 72 h. The ethanol 99.9%, 100 mL, was added to the precipitated product. To eliminate CTAB, the mixture was sonicated for 30 min at 60 °C [44] and then washed with DI water several times and dried at 80 °C for 24 h.

The experimental procedure was repeated more than three times to verify that the results are reproducible.

2.2.2. Characterization

The mineral compositions in BA and MMS were determined by X-ray fluorescence (XRF), (S2175 Ranger, Bruker, Burladingen, Germany). Fourier transform infrared spectroscopy FT-IR (Perkin Elmer Model Spectrum GX) was used to analyze the specific functional groups of the adsorbents by compressed samples into KBr pellets and then analyzed with a Nicole IS10 spectrometer over the wavelength ranged from 400 to 4000 cm⁻¹. X-ray diffraction (XRD) pattern (PAN analytical, X’Pert Pro MPD) was performed on a Bruker AXS Advance instrument for confirmation the structure. The surface morphology (SEM) of magnetic mesoporous silica was examined by scanning electron microscopy (SEM, Apreo, FEI, South Moravian Region, Czech Republic) running by Schottky field emission at the accelerating voltage of 20 kV in high vacuum mode. The elemental mapping analysis of the sample was recorded by energy-dispersive X-ray spectrometer (EDS)—(X-Max80, Oxford, UK).

2.2.3. Adsorption Experiments

MMS was used for tetracyclines (TC) adsorption in an aqueous solution. To find the ideal conditions, the impact of adsorption time and temperature on the unit adsorption capacity and adsorption rate of TC was investigated. MMS (3 mg) was dispersed in a flask containing 10.0 mL TC solution with various concentrations (10–100 mg/L), pH 6–8, and then fully homogenized with a vortex mixer. The suspension was incubated overnight at 25 °C, 45 °C, and 60 °C and covered with aluminum foil to protect TC from the potential photo degradation. MMS was then separated from the samples through a magnet after centrifuging at 2000 rpm for 15, 30, 45, 60, and 90 min while maintaining the experiment temperature including centrifugation. The residual concentration of TC in an aqueous solution was determined by UV–vis absorbance at 357 nm, using a calibration curve built. All the experiments were replicated thrice, and the averaged results were reported. Equation (1) was used to calculate the percentage removal of TC, and Equation (2) was used to determine the adsorption capacity:

$$Removal\left(\%\right)=\frac{(C_o-C_e) }{C_o}\times 100$$

(1)

$$Q_e=\frac{(C_o-C_e)\times V}{m}$$

(2)

where:

• Q_e is the amount of TC adsorption per unit weight of adsorbent (mg/g);
• C_o is the initial concentration of TCs (mg/mL);
• C_e (mg/L) is the equilibrium concentration of TC;
• V is the solution volume (mL);
• m is the mass of adsorbent (g).

3. Results and Discussion

3.1. Characterization of MMS

The FT-IR peaks of the MMS before and after adsorption of TC were collected in the range from 400 to 4000 cm⁻¹ (Figure 1), indicating the chemical bonds and functional groups in the compound. The O–H stretching vibration was identified as the source of the big broadband at 3440 to 3443 cm⁻¹. The absorption peaks at 1640 cm⁻¹ were caused by the symmetric and asymmetric bending vibration of C=O. Fe–O stretching was assigned to the band below 700 cm⁻¹. The characteristic absorption bands of the sample at 692 and 583 cm⁻¹ were assigned to iron(II) oxide bending, and the band at 459 to 446 cm⁻¹ was ascribed to the bending vibration mode of Fe₂O₃. The band at 1048 cm⁻¹ to 1051 cm⁻¹, which is associated with Si-O-Si antisymmetric stretching vibrations, is a sign that silicon dioxide is present in the sample. The bands at 579 and 583 cm⁻¹ are also an indication of the presence of Si–O–Fe [45]. Most of the adsorbent’s peaks remained constant after TC adsorption, revealing that the adsorption procedure did not modify the structure of material. However, several distinctive peaks at 1478, 2852, and 2922 cm⁻¹ were characteristic peaks of the C=C skeleton and C-H stretching vibration of CH₂ and CH₃ induced by aromatic groups of TC [46,47]. The FT-IR spectrum proved that TC was adsorbed onto the adsorbent and the sample’s structure is mostly stable after adsorption [48].

The XRD profile of the MMS in the FT-IR spectrum proved that TC was adsorbed onto the adsorbent and the sample’s structure is mostly stable after adsorption [48].

Comparisons with those of SiO₂ (extracted from biomass power plant Ash) and Fe₂O₃ peak [49] are shown in Figure 2. The XRD pattern of the MMS exhibited the characteristic diffraction peaks of Fe₂O₃ with weak intensity due to the lower concentration of Fe₂O₃. In addition, the relatively slightly broad peak observed at 15−30° arose from the SiO₂, demonstrating that the crystalline structure of the iron oxide was retained after the encapsulation in silica.

The morphological structures of the MMS were examined using SEM. As shown in Figure 3a, there are many unobvious spherical particles (aggregated) with a size of around 50−80 nm. A rough surface, which mainly consists of not−well−crystallized Fe₂O₃, seems not to be efficient to achieve homogeneous coating SiO₂, but instead irregular and severely agglomerated particles are observed in Figure 3b,c.

The samples typically determined the elemental analysis or chemical properties using energy-dispersive X−ray spectrometry (EDS). The various elements in the sample are represented by energy peaks. As-prepared nanocomposites were found to include significant amounts of Fe, O, Si, Ti, Na, Ca, Al, etc. (Figure 3d), confirming the formation of additional nanocrystals on the surface of Fe₂O₃@SiO₂ particles.

3.2. Adsorption Isotherms of TC on MMS

In order to elucidate the interactions between the adsorption ability of the MMS and TC in solution at adsorption equilibrium, the data were modeled using four adsorption isotherms, including Langmuir, Freundlich, Temkin, and Sips isotherm models at 25, 45, and 60 °C. The isotherm models were listed below with Equations (3)–(6) as follows.

$$Q_e=\frac{Q_m{K}_L{C}_e}{1+K_L{C}_e};R_L=\frac{1}{1+K_o{C}_o}$$

(3)

$$Q_e=K_f{C}_e^{\frac{1}{nf}}$$

(4)

$$Q_e=\frac{RT}{b_T}ln(K_T{C}_e)$$

(5)

$$Q_e=Q_m\frac{{(K_S{C}_e)}^a}{1+{(K_S{C}_e)}^a}$$

(6)

where Q_m = maximum adsorption ability (mg/g), C_e = equilibrium concentration (mg/L), Q_e = adsorption capacity (mg/g) at equilibrium time, K_L = Langmuir constant, R_L = separation constant, K_f and n = Freundlich constant, R = universal gas constant (8.314 J/mol), T = temperature in terms of Kelvin, b_T = Temkin constant, K_T = equilibrium bond constant related to the maximum energy of bond, and K_s and $a$ are Sips constants.

According to Figure 4a, as the equilibrium TC concentration climbed, the TC’s capacity to adsorb on the MMS also dramatically increased. It was determined that the adsorption process was endothermic because the adsorption capacity of TC increased as the adsorption temperature rose. This is likely because the higher temperature supported the quantity of activated functional groups on the surface of the MMS or sped up the diffusion rate of the tetracycline molecules.

The starting TC concentration used in the modeling procedure varied from 10 to 100 mg/L, and the adsorption temperature was adjusted at 25, 45, and 60 °C. Figure 4b–d show the plot of the nonlinear fits for the aforementioned four isotherm models, and

Table 1. Parameters of adsorption isotherm model of TC onto MMS at 25, 45, 60 °C.

Temp (°C)	25 °C	45 °C	60 °C
Langmuir
Qm	18.11 ± 2.65	43.92 ± 15.67	276.74 ± 159.62
KL	0.031 ± 0.009	0.009 ± 0.004	0.002 ± 0.001
R2	0.9832	0.9876	0.9815
SSE	1.58	1.74	0.19
χ2	0.53	0.53	0.53
Freundlich
n	0.54 ± 0.003	0.79 ± 0.001	0.96 ± 3.25 × 10−4
KF	1.31 ± 0.01	0.64 ± 0.003	0.56 ± 6.20 × 10−4
R2	0.9867	0.9983	0.9999
SSE	138.27	32.07	1.66
χ2	0.14	0.03	0.002
Temkin
AT	0.58 ± 0.01	0.58 ± 0.01	0.88 ± 0.03
BT	790.61 ± 7.53	790.62 ± 7.53	909.52 ± 10.28
R2	0.9181	0.9181	0.8880
SSE	855.40	3806.77	8002.25
χ2	0.86	3.89	8.03
Sips
Qm	12.92 ± 0.11	14.52 ± 0.17	17.07 ± 0.24
KS	0.18 ± 0.02	0.35 ± 0.08	0.47 ± 0.13
A	2.09 ± 0.18	3.79 ± 0.77	3.63 ± 0.87
R2	0.9603	0.9415	0.9218
SSE	512.81	1055.65	2210.77
χ2	0.51	1.06	2.22

lists the related parameter values. The Langmuir model was more appropriate to explain the adsorption of TC onto the MMS, as shown by the correlation coefficients, R², which were consistently close to 1 and the sum of square error (SSE), chi-square (χ²) values, which suggest that the surface of the MMS was homogeneous and the adsorption was monolayer. Additionally, K_L is a significant evaluation factor in relation to the binding site affinities. The value of K_L was between 0–1 and declined with the rising temperature, and separation factor (R_L) <1 reflected that the adsorption of TC onto the MMS was favorable [50]. The nonlinear fitting of the Langmuir model (Figure 4b) showed the adsorption capacity (Q_m) of MMS for TC was 276.74 mg/g.

The R² values of the Freundlich isotherms were from 0.98 to 0.99, but the 1/n values were more than 1, at 1.85, 1.26, and 1.04, with increased temperatures at 25, 45, 60 °C, respectively, suggesting the adsorption is not prone to occur. Therefore, the Freundlich isotherm had a hard time accurately explaining the TC adsorption pattern.

The Temkin and Sips isotherms were a poor fit, although heterogeneity was also assumed. The Temkin isotherms illustrate that the heat of adsorption is negatively proportional to the surface area covered by the adsorbate molecules. The Temkin’s constant, abbreviated B_T, represents the heat generated during adsorption. The fact that B_T was positive (790.61–909.52 J/mol) for the TC adsorption from the aqueous solution indicates that TC was endothermically adsorbed on the MMS. The equilibrium binding constant A_T is the quantity that relates to the greatest binding energy. When the temperature was raised from 25 to 60 °C, the values of A_T also increased from 0.58 to 0.88 L/mg, indicating that the adsorption of TC on the MMS was endothermic.

Equations (7) and (8) of the van’t Hoff equation were used to compute the Gibbs free-energy change ΔG° enthalpy ΔH° and entropy ΔS° during the adsorption process (8). The results are displayed in

Table 2. Thermodynamic parameters of TC adsorption onto MMS.

Temperature (K)	ΔG° (kJ/mol)	ΔH° (kJ/mol)	ΔS° (J/(mol.K)
298	−23.61	−7.62	−15.94
318	−21.92
333	−18.79

.

$$\mathsf{\Delta }{G}^{\mathrm{o}}=-\mathit{RT}\mathit{ln}{K}_L$$

(7)

$$ln\left(K_L\right)=\frac{\Delta S^{\mathrm{o}}}{R}-\frac{\Delta H^{\mathrm{o}}}{RT}$$

(8)

where R is the ideal gas constant 8.314 J/(mol·K), T is Kelvin temperature (K), and K_L is the Langmuir isotherm equilibrium constant (L/mg) [50,51,52,53].

The absolute values of ΔG° are negative for all the parameter intervals, indicating that the adsorption behavior is spontaneous. ΔG° also increased as the adsorption temperature rose, suggesting that the adsorption process was impeded [51]. Moreover, the value of ΔG° in the range from −20 to −80 (kJ/mol) showed that physisorption was involved in the process, but a minor chemical action effect could speed up the procedure. Furthermore, the value of ΔH° lower than 20 (kJ/mol) revealed the physical adsorption with van der Waals interaction implied in the process mechanism [52]. A very low ∆S° value (−15.94 J/(mol.K)) proved a little change in entropy occurred during the adsorption of TC by the MMS [53].

3.3. Adsorption Kinetics of TC on MMS

The adsorption kinetics explained the impact of time and temperature on the TC adsorption rate. The adsorption kinetics test was conducted by adding 0.03 g of adsorbent to 10 mL of 100 mg/L TC solution at 45 °C. The adsorption data were fitted using three popular kinetic models, including the pseudo-first-order, pseudo-second-order, and Elovich models. The kinetic models are listed in the Equations (9)–(11) respectively, the corresponding nonlinear curves are shown in Figure 5, and

Table 3. Adsorption Kinetic Coefficients.

Model	Parameter	Value
pseudo-1st-order	K1	0.14 ± 0.02
	Qe	1.44 ± 0.02
	R2	0.9971
	SSE	0.005
	χ2	0.00
pseudo-2nd-order	Qe	1.56 ± 0.003
	K2	0.14 ± 0.002
	R2	0.9646
	SSE	2.44
	χ2	0.00
Elovich	β	1.33 ± 0.05
	α	3.90 ± 0.03
	R2	0.9559
	SSE	4.20
	χ2	4.20

contains the estimated and shown kinetic parameters.

$$Q_t=Q_e(1-e^{-K_{1}t})$$

(9)

$$Q_t=\frac{K_e{Q}_e^{2}t}{1+K_2{Q}_{e}t}$$

(10)

$$Q_t=\left(\frac{1}{b}\right)ln\left(ab\right)+\left(\frac{1}{b}\right)ln\left(t\right);$$

(11)

For the pseudo-second-order and Elovich models, R² was considerably less than 1, and SSE, χ² were higher. It was determined that neither model could adequately explain the experimental findings. However, the pseudo-first-order kinetic model showed a high value of R² (0.9971) and a small value of SSE, χ².The closer fitting curves imply that the pseudo-first-order adsorption process was dominant for TC adsorption on the MMS.

As presented in

Table 4. The capacity adsorption of TC with other different adsorbents’ utilization from waste.

Adsorbent	Qm (mg g−1)	Refs.
Alkali modified magnetic biochar (MSBC)	97.962	[54]
Alkali-acid modified magnetic biochar (MSABC)	98.334
The raw biochar (RBC)	37.803
Alfalfa-derived biochar	372	[55]
Pinus taeda-derived activated biochar	274.8	[56]
Waste chicken-feather-derived multilayered graphene-phase biochar	388.33	[57]
Clay-biochar composites	77.962	[58]
Spent coffee-ground-derived biochar	39.22	[59]
Biomass ash pyrolyzed from municipal sludge	50.75	[60]
Shrimp Shell Waste	229.98	[61]
Magnetic mesoporous silica from BA-Biomass Power Plant	276.74	In this study

, the MMS has a high adsorption capacity when compared to the Q_m values of a variety of adsorbents utilized from waste for TC antibiotics in aquatic environments. From this table, the MMS can be considered as an alternative adsorbent material for TC adsorption.

4. Conclusions

Magnetic mesoporous silica that was fabricated from rubber biomass power plant ash was considered as a low-price adsorbent for tetracycline adsorption in aqueous solution. The adsorption of tetracycline onto the magnetic mesoporous silica was a monolayer endothermic process. The maximum tetracycline adsorption capacity of magnetic mesoporous silica in aqueous solutions was 276.74 mg/g at 60 °C. With the adsorption, the pseudo-first-order model accurately reflects the data on adsorption kinetics, while the Langmuir model for adsorption matches the data on adsorption isotherms well. Tetracycline adsorption by the MMS was spontaneous due to the negative values of the Gibbs free-energy and enthalpy change, and at the higher temperature the adsorption was adversely affected. The adsorbent showed exceptional properties, including tetracycline removal efficiency and easy separation from aqueous media by a magnet. Therefore, magnetic mesoporous silica could be used as a potential adsorbent for tetracycline from an aqueous solution. Furthermore, this adsorbent can be utilized as an environmentally friendly adsorbent for the treatment of wastewater.