Mechanism of Enhanced Fluoride Adsorption Using Amino-Functionalized Aluminum-Based Metal–Organic Frameworks

Abstract

Due to the increasing fluoride concentrations in water bodies, significant environmental concerns have arisen. This study focuses on aluminum-based materials with a high affinity for fluorine, specifically enhancing metal–organic frameworks (MOFs) with amino groups to improve their adsorption and defluorination performance. We systematically investigate the factors influencing and mechanisms governing the adsorption and defluorination behavior of amino-functionalized aluminum-based MOF materials in aqueous environments. An SEM, XRD, and FT-IR characterization confirms the successful preparation of NH₂-MIL-101 (Al). In a 10 mg/L fluoride ion solution at pH 7.0, fluoride ion removal efficiency increases with the dosage of NH₂-MIL-101 (Al), although the marginal improvement decreases beyond 0.015 g/L. Under identical conditions, the fluoride adsorption capacity of NH₂-MIL-101 (Al) is seven times greater than that of NH₂-MIL-101 (Fe). NH₂-MIL-101 (Al) demonstrates effective fluoride ion adsorption across a broad pH range, with superior fluoride uptake in acidic conditions. At a fluoride ion concentration of 7 mg/L, with 0.015 g of NH₂-MIL-101 (Al) at pH 3.0, adsorption equilibrium is achieved within 60 min, with a capacity of 31.2 mg/g. An analysis using adsorption isotherm models reveals that the fluoride ion adsorption on NH₂-MIL-101 (Al) follows a monolayer adsorption model, while kinetic studies indicate that the predominant adsorption mechanism is chemical adsorption. This research provides a scientific basis for the advanced treatment of fluoride-containing wastewater, offering significant theoretical and practical contributions.

1. Introduction

Fluoride is ubiquitously present in nature, acting both as a toxicological marker for water bodies and as an essential trace element for human health and development [1]. In industrial processes, such as those in the solar cell, phosphate fertilizer, cement production, and thermal power generation industries, the byproducts often include fluoride-containing “three wastes”, which contribute to high concentrations of fluoride in wastewater [2]. Additionally, significant amounts of fluoride are released into the atmosphere during volcanic eruptions, where they subsequently settle into water bodies [3,4]. Fluorinated minerals also release fluoride through surface runoff erosion and natural weathering, leading to its accumulation in soil and water bodies [5,6,7]. Research indicates that the primary route for fluoride entry into the human body is through drinking water [8]. According to China’s drinking water hygiene standards (GB5749-2006) [9], fluoride levels in drinking water should be between 0.5 and 1.0 mg/L. This range is critical because insufficient fluoride intake can lead to dental caries, while excessive fluoride can cause fluorosis and skeletal fluorosis, adversely affecting health [10,11]. Concentrations exceeding 10 mg/L can severely impact the environment and human health, disrupting animal and plant reproduction and potentially causing neurological, immune, and reproductive system issues [7,12]. Long-term excessive fluoride intake can lead to memory loss, cognitive decline, and a weakened immune system, increasing susceptibility to diseases [13,14,15,16,17]. Addressing water pollution [18,19,20,21], particularly fluoride contamination, is a pressing global concern. Given the rising fluoride levels in water bodies, it is crucial to develop effective fluoride removal strategies to protect the ecological balance essential for human survival [22].

Common methods for fluoride removal include chemical precipitation, ion exchange, coagulation sedimentation, and adsorption [23,24,25]. Chemical precipitation involves solid–liquid separation by generating precipitates through chemical reactions. These precipitates often require further treatment; if not managed properly, the pollutants in the precipitates may be released back into the environment, causing secondary pollution [26,27]. Moreover, the efficiency of chemical precipitation is influenced by various operating conditions, such as pH, temperature, and stirring speed. Minor variations in these parameters can significantly affect the formation and separation of precipitates [28]. The ion exchange method is a separation technique based on reversible chemical reactions between active groups on ion exchange materials (e.g., ion exchange resins) and ions in the solution. This method effectively removes impurity ions from water and isolates valuable compounds [29,30,31]. However, it is costly, and its regeneration process typically requires substantial amounts of salts to supply exchange ions, leading to high salt consumption [32]. This consumption becomes particularly pronounced in wastewater with high ion concentrations, thereby increasing treatment costs [33]. Coagulation sedimentation is a widely used combined physical–chemical treatment technology in wastewater management. It operates on the principle that colloids and fine suspended particles aggregate into larger flocs under the influence of coagulants. These flocs then settle under gravity, achieving solid–liquid separation and removing harmful substances. This method is effective and versatile, finding applications in domestic, industrial, and rural sewage treatment. However, there is a risk of secondary pollution if coagulants are improperly selected or excessively used, potentially increasing certain wastewater parameters (e.g., pH, color), which poses a risk of secondary pollution. The adsorption method for fluoride removal utilizes specific adsorbents, such as activated alumina, polyaluminum salts, and lignite adsorbents. These materials possess porous structures and large specific surface areas, enabling effective fluoride ion adsorption [34]. When fluoride containing wastewater passes through these adsorbents, fluoride ions are adsorbed on the surface of the adsorbents to promote their removal [35]. The key to improving this method is identifying an adsorbent that offers high effectiveness at low cost [36,37]. An ideal fluoride adsorbent should feature rapid adsorption kinetics, cost-effectiveness, recyclability, high adsorption capacity, and minimal reactivity with other substances [38]. High-performance adsorbents typically have significant porosity and a specific surface area or possess chemical groups that readily react with pollutants [39,40].

The adsorption method distinguishes itself among various water purification technologies due to its superior defluorination performance, low cost, and straightforward manufacturing process, making it the prevalent choice for defluorination at present [41,42]. This method is not only efficient but also cost-effective and highly practical [43,44,45]. Metal–organic frameworks (MOFs) are a novel class of nanomaterials, typically formed through the self-assembly of metal ions coordinated with organic ligands [46,47,48,49]. MOFs are characterized by their high porosity and robust chemical stability, rendering them suitable for applications in catalysis [50,51,52], adsorption and separation [53], optical materials [54], and magnetic materials. These materials exhibit a significant adsorption capacity for specific organic pollutants in aqueous environments, thus offering promising advancements in water purification [55,56]. As a new type of functional molecular material, MOFs have the following advantages compared to activated carbon and zeolite materials: a large specific surface area, adjustable pore size, an ordered microporous structure, diverse pore size and skeleton structure, modifiable functional groups on the pore surface, and unsaturated metal coordination sites [57,58,59]. MOFs can be categorized into various types, such as Zeolitic Imidazolate Frameworks (ZIFs), Materials of Institut Lavoisier (MILs), and University of Oslo (UIO) series, based on their metal ions and organic ligands [60]. Current synthesis methods for MOFs include microwave-assisted synthesis, ultrasonic methods, liquid-phase diffusion, solvothermal synthesis, and mechanical stirring [61]. To enhance the porosity and performance of MOFs, several activation techniques are employed, including vacuum heat treatment, solvent exchange, supercritical carbon dioxide treatment (scCO₂), freeze-drying, ultrasonic treatment, and chemical activation [62]. These activation methods effectively improve MOF porosity, thereby optimizing their performance in catalysis and adsorption applications [53].

This study successfully synthesized NH₂-MIL-101 (Al) material using a solvothermal method and confirmed its synthesis quality through comprehensive material characterization. The fluoride ion removal efficiency and mechanism of this material were assessed using adsorption experiments. A series of controlled comparative experiments were conducted to explore the impacts of varying dosages, initial fluoride ion concentrations, pH levels, reaction times, and temperatures on the fluoride ion removal efficiency of NH₂-MIL-101 (Al). The optimal dosage and conditions for fluoride ion adsorption by NH₂-MIL-101 (Al) were determined, providing a robust theoretical and experimental foundation for its practical applications.

2. Materials and Methods

2.1. Instruments and Experimental Reagents

Sodium fluoride (NaF, analytical grade), 2-aminoterephthalic acid (C₈H₇NO₄, 98%), and acetic acid (CH₃COOH, 99.5%) were obtained from Shanghai McLean Biochemical Technology Co., Ltd. (Shanghai, China). N,N-Dimethylformamide (C₃H₇NO, analytical grade) was sourced from Taicang Hujian Reagent Co., Ltd. (Taicang, China). Sodium chloride (NaCl, superior grade), aluminum chloride hexahydrate (AlCl₃·6H₂O, analytical grade), methanol (CH₄O, analytical grade), sodium hydroxide (NaOH, analytical grade), trisodium citrate (C₆H₅Na₃O₇, analytical grade), and hydrochloric acid (HCl, analytical grade) were purchased from China National Pharmaceutical Group Chemical Reagent Co., Ltd. (Beijing, China).

A digital pH meter (PHS-3E, Shanghai Yidian Scientific Instrument Co., Ltd., Shanghai, China), high-speed desktop centrifuge (TG16-WS, Hunan Xiangyi Centrifuge Instrument Co., Ltd., Changsha, China), electronic balance (FA2004B, Shanghai Youke Instrument Co., Ltd., Shanghai, China), constant-temperature drying oven (202 type, Beijing Yongguangming Medical Instrument Co., Ltd., Beijing, China), rotary vane vacuum pump (2XZ type, Jiaojiang Shuang’e Vacuum Equipment Factory, Taizhou, China), ultrasonic cleaning machine (820HT, Changsha Mingjie Instrument Co., Ltd., Changsha, China), constant-temperature oscillator (XMTE-205, Changzhou Guowang Instrument Co., Ltd., Changzhou, China), and deionized water preparation unit (LAN-19-H, Lide Water Treatment Equipment Co., Ltd., Chongqing, China) were obtained for this study.

2.2. Preparation of NH₂-MIL-101 (Al)

NH₂-MIL-101 (Al) was synthesized using a solvothermal method [63,64]. Precisely weigh 0.51 g of AlCl₃·6H₂O and 0.56 g of 2-aminoterephthalic acid using an analytical balance. Dissolve these reagents in 30 mL of N,N-dimethylformamide. Subject the resulting solution to ultrasonic agitation for 30 min to ensure thorough mixing. Transfer the homogenized solution into a 50 mL reaction vessel, seal it tightly, and perform the reaction at 130 °C for 72 h. Allow the reaction mixture to cool to room temperature naturally. After cooling, separate the product by centrifugation at 11,000 rpm to obtain a yellow precipitate. Wash the precipitate with N,N-dimethylformamide to remove the supernatant, then perform three additional washes with anhydrous methanol. Finally, place the washed precipitate in a vacuum drying oven at 60 °C and dry overnight to obtain NH₂-MIL-101 (Al) as a light-yellow powder. The preparation method of NH₂-MIL-101 (Fe) is based on our previous report [65].

2.3. Characterization Methods for NH₂-MIL-101 (Al)

(1)

X-ray Diffraction Analysis (XRD)

To investigate the crystal structure and phase composition of NH₂-MIL-101 (Al), an X-ray diffraction analysis was performed using a MiniFlex600 X-ray powder diffractometer (Rigaku Co., Ltd., Tokyo, Japan). Copper Kα radiation was utilized as the X-ray source to ensure high precision in the measurements. The diffraction measurements were conducted at a current of 40 mA and a voltage of 40 kV to maintain a stable instrument operation. The data were collected with a scanning speed of 10°/min over a 2θ range of 5° to 90° to capture comprehensive diffraction information of the material.

(2)

Fourier Transform Infrared Spectroscopy (FT-IR)

Fourier Transform Infrared Spectroscopy (FT-IR) was employed to analyze the chemical composition, functional groups, and intermolecular interactions of NH₂-MIL-101 (Al). This technique utilizes the absorption and transmission characteristics of infrared light to generate spectra that reveal detailed information about chemical bonds and functional groups within the material. The FT-IR spectra were obtained using a Nicolet iS50 spectrometer (Thermo Fisher Scientific, Inc., Waltham, MA, USA), with samples prepared using the KBr pellet method, covering a wavelength range of 400 to 4000 cm⁻¹. These spectral data provide valuable insights into the chemical properties of NH₂-MIL-101 (Al).

(3)

Scanning Electron Microscopy (SEM)

To examine the surface microstructure, elemental distribution, and particle size of NH₂-MIL-101 (Al), a JSM-7610FPlus scanning electron microscope (JEOL Co., Ltd., Tokyo, Japan) was employed. This high-resolution microscopy technique allowed for detailed observation of the fine surface structure and precise measurement of particle size.

(4)

Specific Surface Area and Pore Size Analysis (BET)

The specific surface area and pore structure of NH₂-MIL-101 (Al) were analyzed using the Brunauer–Emmett–Teller (BET) method. A JWGB JY-BK 112 specific surface area analyzer (Beijing JWGB Sci.& Tech. Co., Ltd., Beijing, China) was utilized to determine key parameters such as the specific surface area, pore volume, and pore size of the sample. This analysis not only enhances the accuracy of the research but also provides a deeper understanding of the microstructural characteristics of NH₂-MIL-101 (Al).

2.4. Adsorption Experiment and Analysis

2.4.1. Adsorption Experiment

Fluoride ion concentrations were determined using the standard curve method, employing a pH meter equipped with a fluoride ion selective electrode [66]. The electromotive force (E), measured in millivolts by the pH meter, decreases as the fluoride ion concentration in the solution increases. Before each experimental batch, it is essential to verify the electrode’s actual slope and re-construct the standard curve. To assess the impact of the material dosage and initial fluoride ion concentration on the fluoride ion removal efficiency of NH₂-MIL-101 (Al), the following experiments were conducted:

Material Dosage Effect: Add 0.005 g, 0.010 g, 0.015 g, and 0.020 g NH₂-MIL-101 (Al) to 150 mL of a 10 mg/L fluoride ion solution, adjusting the pH to 7.0–7.5. Then, keep the reaction solution in a constant temperature stirrer at 25 °C for 120 min. After the reaction, filter the solution using a 0.22 µm organic filter, measure the equilibrium concentration, and calculate the equilibrium adsorption capacity ($q_e$).

Initial Concentration Effect: Prepare fluoride ion solutions at concentrations of 1, 2, 3, 4, 5, 6, 7, 8, 9, and 10 mg/L. Add 150 mL of each fluoride ion solution to separate beakers and introduce 0.015 g of NH₂-MIL-101 (Al). React for 120 min, then filter the solutions using a 0.22 µm organic filter. Measure the equilibrium concentrations and determine the equilibrium adsorption capacity ($q_e$).

To evaluate the influence of solution pH and reaction time on the fluoride ion removal performance of NH₂-MIL-101 (Al), the following experiments were performed:

Reaction Time Effect: Prepare 150 mL of a 7 mg/L fluoride ion solution with the pH adjusted to 7, and add 0.015 g of NH₂-MIL-101 (Al). React at 25 °C in a constant temperature mixer, and filter the sample solution every 10 min. Measure the equilibrium concentration after each time point and calculate the adsorption capacity ($q_e$).

pH Effect: Prepare 150 mL of a 7 mg/L fluoride ion solution and sequentially adjust the pH to 1, 3, 5, 7, and 9. Add 0.015 g of NH₂-MIL-101 (Al) to each solution and react at 25 °C for 60 min in a constant temperature stirrer. After the reaction, filter the solution using a 0.22 µm organic filter, measure the equilibrium concentrations, and calculate the equilibrium adsorption capacity ($q_e$).

2.4.2. Adsorption Kinetics

Two kinetic models—the pseudo-first-order kinetic model and the pseudo-second-order kinetic model—were utilized to analyze and describe the interaction between the adsorbent and the adsorbate under specific temperature conditions [67,68].

(1)

Pseudo-First-Order Kinetic Model

The pseudo-first-order rate equation, based on solid adsorption capacity, is frequently employed. If kinetic data fit the pseudo-first-order model, it indicates that diffusion is the predominant mechanism in the adsorption process. The linear equation for this model is given by the following:

$$\ln (q_e-q_t)=\ln q_e-k_{1}t$$

(1)

where t is the adsorption time (minutes), q_t is the adsorption capacity at time t (mg/g), q_e is the adsorption capacity at equilibrium (mg/g), and k₁ is the pseudo-first-order rate constant (min⁻¹) [41].

(2)

Pseudo-Second-Order Kinetic Model

Assuming that the adsorption rate is governed primarily by a chemical adsorption mechanism, the pseudo-second-order kinetic model can be applied. This model suggests that the adsorption involves electron transfer between the adsorbent and the adsorbate. If the kinetic data align well with the pseudo-second-order model, it implies that the adsorption process is mainly controlled by chemical interactions. The linear equation is expressed as the following:

$${\displaystyle \frac{t}{q_t}}={\displaystyle \frac{1}{k_2{q}_e^2}}+{\displaystyle \frac{t}{q_e}}$$

(2)

where t is the adsorption time (min), q_t is the adsorption capacity at time t (mg/g), q_e is the adsorption capacity at equilibrium (mg/g), and k₂ is the pseudo-second-order rate constant (g/(mg·min)).

2.4.3. Adsorption Isotherms

Two widely accepted adsorption models—the Langmuir adsorption isotherm model and the Freundlich adsorption isotherm model—are employed for detailed simulation and analysis [69,70].

(1)

Langmuir Adsorption Isotherm Model

The Langmuir adsorption isotherm model presumes a constant number of adsorption sites on the adsorbent, each with identical adsorption potential. Once a solute molecule occupies an adsorption site, it remains fixed and does not migrate to other sites, with no mutual influence between adsorbed molecules. This model provides a theoretical foundation for understanding the isothermal adsorption process. The linear form of this model is as follows:

$${\displaystyle \frac{1}{q_e}}={\displaystyle \frac{1}{K_L{q}_{max}}}{\displaystyle \frac{1}{C_e}}+{\displaystyle \frac{1}{q_{max}}}$$

(3)

where q_e is the equilibrium adsorption capacity (mg/g), K_L is the adsorption constant, q_max is the theoretical maximum adsorption capacity (mg/g), and C_e is the equilibrium concentration (mg/L).

The adsorption constant (K_L) reflects the adsorption difficulty: a K_L value of 0 indicates an irreversible reaction; K_L between 0 and 1 signifies a feasible reaction; K_L equal to 1 represents linear adsorption; and K_L greater than 1 suggests a more challenging reaction.

(2)

Freundlich Adsorption Isotherm Model

The Freundlich adsorption isotherm model describes adsorption as a non-uniform distribution across multiple molecular layers. This model indicates that adsorption is not evenly distributed but occurs in a multi-layer structure with varying adsorption capacities at different sites. The model suggests that as the adsorbate concentration increases, so does the adsorption, highlighting the site-specific adsorption capacities and the complexity of multi-layer adsorption. The linear form is the following:

$$ln{q}_e={\displaystyle \frac{1}{n}}ln{C}_e+ln{K}_F$$

(4)

where q_e is the equilibrium adsorption capacity (mg/g), and K_F is the adsorption constant.

For plotting isothermal adsorption data, lnq_e is typically plotted against lnC_e. A straight line is fitted to the experimental data, from which the Freundlich parameters K_F and n are derived. K_F is directly related to adsorption capacity, a higher K_F indicates stronger adsorption. The n value represents the adsorption intensity; n > 1 usually indicates that the adsorption process is relatively favorable.

3. Results and Discussion

3.1. Characterization Analysis of NH₂-MIL-101 (Al)

3.1.1. X-ray Diffraction Analysis (XRD)

Figure 1 presents the XRD pattern of NH₂-MIL-101 (Al). The data reveal that the diffraction peaks exhibit high intensity. A comparison with the literature data indicate that the characteristic diffraction peaks of the material exhibit high diffraction intensity and align with the peaks documented in the literature [71,72]. The absence of impurity peaks confirms the successful synthesis of NH₂-MIL-101 (Al). These results suggest that the synthesized NH₂-MIL-101 (Al) material possesses a crystal phase structure consistent with that described in the literature, thereby validating its identity as NH₂-MIL-101 (Al).

3.1.2. Fourier Transform Infrared Spectroscopy (FT-IR)

The chemical bonds in NH₂-MIL-101 (Al) were analyzed using Fourier Transform Infrared Spectroscopy (FT-IR), as illustrated in Figure 2. According to reference [43], the absorption peak at 3491.9 cm⁻¹ corresponds to the asymmetric and symmetric stretching vibrations of the -NH₂ group in the organic ligand. The characteristic peaks at 1569.1 cm⁻¹ and 1442.1 cm⁻¹ are attributed to the N-H bending vibration and C-H stretching vibrations in aromatic rings, respectively. These results confirm the successful synthesis of the NH₂-MIL-101 (Al) material.

3.1.3. Scanning Electron Microscopy (SEM)

Figure 3 reveals that the NH₂-MIL-101 (Al) material exhibits a tightly packed octahedral grain arrangement, forming clusters with a rough surface, which enhances adsorption capacity. This observation is consistent with the findings reported in the literature [73].

3.1.4. Specific Surface Area and Pore Size Analysis (BET)

The specific surface area and pore size distribution of NH₂-MIL-101 (Al) are presented in Figure 4. Figure 4a shows the N₂ adsorption isotherm at 77 K. The low-pressure region indicates a low adsorption capacity for N₂, while the capacity significantly increases as the relative pressure approaches 1.0, exhibiting a Class II adsorption isotherm with an overall S-shape [42]. Figure 4b illustrates the pore size distribution of NH₂-MIL-101 (Al), with a concentration around 20 nm and an average pore size of 17.19 nm. The material is categorized as mesoporous, with an average pore volume of 0.28 cm³/g and a BET specific surface area of 676.94 m²/g.

3.2. Effect of Material Dosage on Adsorption Performance

Figure 5 illustrates the impact of different dosages on fluoride ion removal by NH₂-MIL-101 (Al). Increasing the dosage enhances the adsorption capacity of NH₂-MIL-101 (Al) for fluoride ions. At dosages of 0.005 g and 0.010 g, the equilibrium adsorption capacities are 25 mg/g and 30 mg/g, respectively. Compared to 0.005 g and 0.010 g, he adsorption capacity increases to 36.6 mg/g at 0.015 g, the adsorption capacity increases to 36.6 mg/g, and the removal rate rising from 34% and 40% to 48%, respectively. At a dosage of 0.020 g, the equilibrium adsorption capacity reaches 38.4 mg/g, and the removal rate improves from 48% to 51%. This change indicates that beyond 0.015 g, further increases in dosage do not significantly enhance adsorption performance. With the increase in adsorbent dosage (under the condition of constant fluoride solution concentration and volume), the removal rate increases, which may be due to the increase in the number of active sites on the MOF and the increase in surface area in contact with the solution. However, as the amount of MOFs increases, the increase in removal rate is not significant, because the number of remaining adsorption sites increases during the removal process [63]. Subsequent experiments were therefore conducted with 0.015 g of NH₂-MIL-101 (Al).

Under identical conditions, a 120 min reaction time achieved equilibrium adsorption of fluoride. Figure 5 shows that NH₂-MIL-101 (Al) exhibits a markedly superior adsorption capacity for fluoride compared to NH2-MIL-101 (Fe). Specifically, NH₂-MIL-101 (Al) achieves an adsorption capacity of 36.6 mg/g, whereas NH2-MIL-101 (Fe) only reaches 5.9 mg/g. According to reference [74], this disparity is attributed to NH₂-MIL-101 (Al)’s mesoporous structure, larger specific surface area, and abundant functional groups, which confer a substantial advantage in fluoride ion adsorption over NH₂-MIL-101 (Fe) and enhance the interaction between NH₂-MIL-101 (Al) and fluoride ions.

3.3. Effect of Initial Fluoride Solution Concentration on the Defluorination Efficiency of NH₂-MIL-101 (Al)

Figure 6 depicts the relationship between the equilibrium adsorption capacity (q_e) of NH₂-MIL-101 (Al) and the initial concentration (C₀) of fluoride ion solution. The data demonstrate that as the fluoride ion concentration increases, the adsorption capacity of NH₂-MIL-101 (Al) also rises until it reaches an equilibrium state where further increases in concentration no longer affect the adsorption capacity.

The underlying mechanism can be explained as follows [75]: at lower fluoride ion concentrations, the number of available adsorption sites on NH₂-MIL-101 (Al) significantly exceeds the number of free fluoride ions in the solution, resulting in incomplete adsorption equilibrium. As the concentration of fluoride ions increases, the number of free fluoride ions also increases, providing more ions for adsorption onto the NH₂-MIL-101 (Al) crystals. This increase enhances the adsorption capacity by facilitating the movement of fluoride ions to the available adsorption sites. However, when the initial fluoride ion concentration reaches 7.0 mg/L, the adsorption capacity begins to stabilize. This stabilization is likely due to the saturation of adsorption sites on NH₂-MIL-101 (Al), where additional fluoride ions can no longer be absorbed effectively.

In summary, when the fluoride ion concentration is below 7.0 mg/L, increasing the initial concentration enhances the adsorption performance of NH₂-MIL-101 (Al). Once the concentration reaches 7.0 mg/L, the adsorption capacity remains relatively constant. As shown in Figure 6, at a fluoride ion concentration of 7.0 mg/L, the equilibrium adsorption capacity is 35.1 mg/g and does not increase further with higher initial fluoride ion concentrations. Therefore, subsequent experiments will utilize a fluoride ion solution with a concentration of 7 mg/L.

3.4. Effect of Solution pH on Fluorination Removal of NH₂-MIL-101 (Al)

From Figure 7, it can be seen that the initial solution pH has little effect on the defluorination of the NH₂-MIL-101 (Al) adsorbent, indicating that NH₂-MIL-101 (Al) still has an excellent defluorination performance in a larger pH range, and its adsorption performance is even better under acidic conditions. Under acidic and neutral conditions, the adsorption capacity of NH₂-MIL-101 (Al) can reach over 25 mg/g, and when the pH reaches 9.0, the adsorption capacity begins to decrease to below 20 mg/g. The possible mechanism is the consumption of hydroxyl ions by hydrogen ions in water, further promoting the interaction between fluoride ions and NH₂-MIL-101 (Al) under acidic conditions [73]. Similarly, there is a strong competitive relationship between fluoride ions and hydroxyl ions under alkaline conditions, which leads to the inability of NH₂-MIL-101 (Al) to fully adsorb free fluoride ions in water, resulting in weaker adsorption performance. In order to make the adsorption effect more obvious in subsequent experiments, the pH was adjusted to three before conducting further experiments.

3.5. Effect of Reaction Time on Defluorination Using NH₂-MIL-101 (Al)

Determining the optimal adsorption time is essential for effective wastewater treatment using NH₂-MIL-101 (Al). A slower adsorption rate extends the time required to achieve equilibrium, which may be suboptimal for practical applications in wastewater treatment. Figure 8 illustrates the relationship between adsorption time (t) and the amount of fluoride ions adsorbed (q_t) by NH₂-MIL-101 (Al). The data indicate that qt increases progressively with an adsorption time up to 60 min. Beyond this period, the adsorption capacity stabilizes with minimal variation, suggesting that NH₂-MIL-101 (Al) has reached its adsorption saturation point for fluoride ions. Based on repeated experiments, this saturation level is considered the equilibrium adsorption capacity, with an average q_t of 31.2 mg/g. Compared with previous studies on metal–organic frameworks, the MOF synthesized in this study has a higher q_e and a shorter time required to reach equilibrium [76,77,78,79]. Compared to other adsorbents investigated for F⁻ removal, NH₂-MIL-101(Al) adsorbs F⁻ in a shorter period of time, confirming the appropriate nature of the studied MOF for the purification of F⁻-polluted water [63,79,80].

3.6. Adsorption Kinetics Analysis

To further elucidate the adsorption kinetics of fluoride ions on NH₂-MIL-101 (Al), both pseudo-first-order and pseudo-second-order kinetic models were applied to analyze the adsorption data.

3.6.1. Pseudo-First-Order Kinetic Model

The pseudo-first-order kinetic model was evaluated by plotting the adsorption time (t) on the x-axis against ln(q_e − q_t) on the y-axis, as shown in

Table 1. Parameters of the pseudo-first-order kinetic model.

${\mathit{q}}_{\mathit{e}}$ (mg·g−1)	${\mathit{k}}_1$ (min−1)	${\mathit{R}}^{\mathbf{2}}$
30.4	−0.02346	0.95278

. The resulting linear equation enabled the determination of the model parameters. By calculating the slope and intercept, the equilibrium adsorption capacity (q_e) and rate constant (k₁) were obtained. An analysis of Figure 9 and Table 1 reveals that the regression coefficient (R²) for the pseudo-first-order model was 0.95278.

3.6.2. Pseudo-Second-Order Kinetic Model

For the pseudo-second-order kinetic model, the adsorption time (t) was plotted against t/q_t, as presented in

Table 2. Parameters of the pseudo-second-order kinetic model.

qe (mg·g−1)	K2 (g/(mg·min))	R2
37.0	0.000373	0.99727

. An analysis of Figure 10 and Table 2 shows that the experimental data aligned closely with the fitted line, yielding a high regression coefficient (R²) of 0.99727.

Considering the R² values of the kinetic equations and comparing the calculated and experimental values of equilibrium capacities, the pseudo-second-order kinetic model is more consistent with the obtained experimental results. Therefore, NH₂-MIL-101 (Al) exhibits strong binding chemical adsorption towards fluoride ions [81]. Studies showed that the adsorption of fluoride ions by Al-MOF is achieved by replacing Al−OH on Al atoms to form a stable Al−F structure [82].

3.7. Adsorption Isotherm Analysis

3.7.1. Langmuir Model

To analyze the Langmuir model adsorption isotherm, plot “1/q_e” on the horizontal axis and “1/C_e” on the vertical axis for three different temperatures, as illustrated in Figure 11. Linear regression of the data yields the fitting equations and parameters listed in

Table 3. Langmuir adsorption isotherm parameters.

Temperature/k	Slope	Intercept	qm	KL	R2
298	0.01501	0.05606	67	0.2677	0.9942
308	0.01265	0.0501	79	0.2525	0.9801
318	0.016	0.02593	63	0.6170	0.9879

.

The data presented in Table 3 indicate that the fluoride ion adsorption by NH₂-MIL-101 (Al) is significantly influenced by temperature. As temperature increases, the adsorption capacity for fluoride ions also rises. Furthermore, the R² values remain above 0.98 across different temperatures, suggesting that the adsorption process conforms well to the Langmuir model. This trend implies that the underlying mechanism for fluoride ion adsorption by NH₂-MIL-101 (Al) is primarily single-molecule adsorption.

3.7.2. Freundlich Model

For the Freundlich model adsorption isotherms, plot “ln(C_e)” on the horizontal axis and “ln(q_e)” on the vertical axis for three different temperatures, as shown in Figure 12. The fitting results and corresponding parameters are summarized in

Table 4. Freundlich adsorption isotherm parameters.

Temperature/k	Slope	Intercept	1/n	KF	R2
298	1.1222	0.6731	0.6731	13.2520	0.9761
308	1.1954	0.6199	0.6199	15.6830	0.9547
318	1.3320	0.5199	0.5199	21.4822	0.9417

.

Figure 12 shows that the Freundlich model fits the NH₂-MIL-101 (Al) fluoride ion adsorption data less effectively compared to the Langmuir model. The n values, which are all below 0.7, suggest that the adsorption process is not easily accomplished. Additionally, the linear fitting coefficients are consistently below 0.98, indicating a poorer fit compared to the Langmuir model.

Comparing the results from the two models, it is evident that the Langmuir model provides a superior fit for fluoride ion adsorption by NH₂-MIL-101 (Al) compared to the Freundlich model. This result confirms that single-molecule adsorption is the dominant mechanism in this process [83,84].

4. Conclusions

This study successfully synthesized NH₂-MIL-101 (Al) material using a solvothermal method and systematically investigated its effectiveness in removing fluoride ions from the solution, along with the underlying mechanisms and influencing factors. Characterization via SEM, XRD, and FT-IR confirms the successful preparation of the NH₂-MIL-101 (Al) material. In a solution containing 10 mg/L fluoride ions at pH 7, fluoride removal efficiency improves with an increasing NH₂-MIL-101 (Al) dosage; however, the marginal benefit decreases when the dosage exceeds 0.015 mg/g. Under identical conditions, the fluoride adsorption capacity of NH₂-MIL-101 (Al) is seven times greater than that of NH₂-MIL-101 (Fe). NH₂-MIL-101 (Al) demonstrates effective fluoride ion adsorption across a broad pH range, with enhanced adsorption in acidic conditions, outperforming NH₂-MIL-101 (Fe). When the fluoride ion concentration is 7 mg/L, and 0.015 g of NH₂-MIL-101 (Al) is used at pH 3.00, the material achieves adsorption equilibrium within 60 min, with an equilibrium capacity of 31.2 mg/g. The analysis using adsorption isotherm models indicates that fluoride ion adsorption onto NH₂-MIL-101 (Al) follows a monolayer adsorption pattern. The kinetic analysis reveals that the predominant mechanism of fluoride ion adsorption is chemical adsorption. The synthesis of NH₂-MIL-101 (Al) has good scalability in industrial scale water treatment applications. However, cost and reuse are the main challenges and considerations for scaling up the process. This research provides a scientific basis for the advanced treatment of fluoride-containing wastewater, offering significant theoretical and practical implications.