Efficient Adsorptive Removal of Phosphonate Antiscalant HEDP by Mg-Al LDH

Abstract

Phosphonate-based antiscalants such as 1-hydroxyethane-1,1-diphosphonic acid (HEDP) are extensively employed in industrial water treatment but pose significant environmental challenges due to their persistence and phosphorus content. In this study, Mg-Al layered double hydroxide (Mg-Al LDH) was synthesized and evaluated for its capacity to adsorb and remove HEDP. Mg-Al LDH showed a pronounced adsorption affinity and an exceptionally high capacity of 276.0 mg g⁻¹ at pH 7.0. The adsorption process was remarkably fast, attaining 97% of equilibrium uptake within 45 min at 298 K. The adsorption data fit well to the Elovich kinetic model and the Langmuir isotherm, indicating that the adsorption process is dominated by chemisorption. Thermodynamic analysis further confirmed its spontaneous nature. Additionally, Mg-Al LDH demonstrated strong tolerance to environmental fluctuations. Characterization techniques, including XRD, FTIR, and zeta potential measurements, confirmed that HEDP adsorption onto Mg-Al LDH primarily occurs via surface complexation with metal sites and electrostatic attraction. These findings demonstrate that Mg-Al LDH is a highly effective adsorbent for removing persistent phosphonate pollutants from wastewater streams.

1. Introduction

1-Hydroxyethane-1,1-diphosphonic acid (HEDP, C₂H₈O₇P₂) is among the most commonly used organophosphonate scale inhibitors in industrial water treatment [1]. Owing to its excellent chelating ability and high chemical stability, it is extensively applied in cooling towers, boilers, and reverse osmosis systems to prevent mineral scaling [2]. However, its environmental persistence and high phosphorus content have raised serious ecological concerns, particularly in accelerating eutrophication and contributing to long-term water pollution [3,4]. Conventional wastewater treatment methods, including biodegradation and chemical precipitation, are often ineffective for removing HEDP, prompting the search for more efficient removal strategies.

Among the available approaches, adsorption has attracted considerable attention for its operational simplicity, cost-effectiveness, and potential for phosphorus resource recovery [5,6,7]. The selection of a phosphorus-locking agent is critical, as it must possess both high adsorption capacity and good stability to form stable phosphate minerals and prevent secondary release. However, current studies reveal that available adsorbents for HEDP still face limitations in performance or cost. For example, Yuan et al. employed lanthanum-loaded biochar for HEDP adsorption, achieving a removal capacity of 15.80 mg g⁻¹ [8]. Li et al. developed magnetic Fe/Eu@PAC composites with a capacity of only 27.45 mg-P g⁻¹ [9]. Wan et al. prepared kaolinite/lanthanum carbonate composites with a theoretical adsorption capacity of just 4.93 mg-P g⁻¹ [10]. Unfortunately, the high cost of rare-earth lanthanum limits their feasibility for large-scale applications. These findings highlight the urgent need for high-performance and inexpensive adsorption materials for removing HEDP.

Layered double hydroxide (LDH) have attracted considerable attention as adsorbents for phosphorus removal, owing to their unique layered architecture, excellent stability, high surface area, and compositional tunability of both metal cations and interlayer anions [11,12,13]. The positively charged metal hydroxide layers, intercalated with exchangeable anions and water molecules, confer a strong affinity toward anionic pollutants, making LDH highly effective in water treatment applications. Zhao et al. reported that Ca-Fe LDH exhibited an exceptionally high adsorption capacity for HEDP (335.7 mg P g⁻¹) and formed a highly ordered Ca-Fe-HEDP ternary composite through a phase transformation process, highlighting the great potential of LDHs for removing phosphorus-containing antiscalants [14]. However, the removal mechanism of Ca-Fe LDH involves significant phase transformation accompanied by the release of Ca²⁺ and Fe³⁺ into the solution, which becomes more pronounced at low HEDP concentrations (<200 mg L⁻¹). In contrast, Mg-Al LDH is regarded as a particularly promising candidate for practical applications due to its low-cost precursors, environmental compatibility, and outstanding structural stability [15,16]. Nevertheless, the potential of Mg-Al LDH for the adsorption and removal of HEDP has not yet been explored.

In this work, two-dimensional Mg-Al LDH was prepared by a co-precipitation method, and its performance in adsorbing HEDP from wastewater was systematically investigated. The microstructure, crystal phase, chemical composition, specific surface area and pore size distribution, and surface potential were characterized using Scanning electron microscopy (SEM), X-ray diffraction (XRD), Fourier-transform infrared (FTIR), Brunauer–Emmett–Teller (BET), and zeta potential measurements, respectively. The influence of pH, coexisting ions, adsorbent dosage, initial HEDP concentration, and temperature on adsorption behavior was investigated. Adsorption kinetics, isotherms, and thermodynamic parameters were analyzed to clarify the adsorption process. In addition, XRD, FTIR, and zeta potential data were used to preliminarily probe the underlying adsorption mechanism.

2. Materials and Methods

2.1. Reagents and Materials

All chemicals used in this study were of analytical grade and used as received without further purification. Magnesium nitrate hexahydrate (Mg(NO₃)₂·6H₂O) and aluminum nitrate nonahydrate (Al(NO₃)₃·9H₂O) were obtained from Aldrich Chemical Co. (Shanghai, China). Sodium hydroxide (NaOH, >96%), sulfuric acid (H₂SO₄, 73–75%), ethanol (EtOH, ≥99.7%), ammonium molybdate tetrahydrate ((NH₄)₆Mo₇O₂₄·4H₂O), and potassium antimony tartrate (C₃H₃KO₇Sb, 99%) were purchased from Sinopharm Chemical Reagent Co., Ltd. (shanghai, China). Ultrapure water (18.2 MΩ·cm) was used throughout the experiments.

2.2. Preparation of Mg-Al LDH

Mg-Al LDH was prepared via a co-precipitation method. Briefly, 75 mM Mg(NO₃)₂·6H₂O and 25 mM Al(NO₃)₃·9H₂O were dissolved in 100 mL of ultrapure water under a nitrogen atmosphere. A 2 M NaOH solution was then added dropwise at a rate of 2 mL min⁻¹ until the pH reached ≥10, with the nitrogen atmosphere maintained throughout the process. The resulting suspension was aged at 60 °C for 12 h, followed by repeated washing with deionized water to remove residual salts. Finally, the product was freeze-dried at −50 °C for 24 h to obtain Mg-Al LDH powder.

2.3. Mg-Al LDH Characterizations

The crystal structure and interlayer spacing of the Mg-Al LDH were investigated by X-ray diffraction (XRD) on a Rigaku Ultima IV diffractometer (Rigaku Corporation, Tokyo, Japan) with Cu Kα radiation (λ = 0.15418 nm) at 40 kV and 40 mA. Scans were performed from 5° to 80° (2θ) at a rate of 5° min⁻¹. The interlayer distance was derived from Bragg’s law (Equation (1)) [17]:

$$2\mathrm{d}\mathrm{s}\mathrm{i}\mathrm{n}\mathrm{\theta }=\mathrm{n}\mathrm{\lambda }$$

(1)

where d represents the interplanar spacing (nm), θ denotes the Bragg angle, n is the diffraction order (n = 1), and λ represents the wavelength of the X-ray (1.5418 Å).

Sample morphology was observed using scanning electron microscopy (SEM, Zeiss Gemini SEM 300, Carl Zeiss AG, Jena, Germany), which utilizes electron beams to produce surface images [18]. Functional groups and chemical bonds were identified via Fourier-transform infrared (FTIR) spectroscopy (Nicolet 380, Thermo Fisher Scientific, Waltam, MA, USA) across 400–4000 cm⁻¹. Elemental composition and oxidation states were examined by X-ray photoelectron spectroscopy (XPS, ESCALAB 250Xi, Thermo Fisher Scientific, Waltham, MA, USA) with an Al Kα source (12.5 kV). Surface area and pore size distribution were determined by the Brunauer–Emmett–Teller (BET) method using a Micromeritics ASAP 2460 analyzer (Micromeritics Instrument Corporation, Norcross, GA, USA), with approximately 0.1 g of sample under N₂ atmosphere.

2.4. HEDP Adsorption Experiment

2.4.1. Influence of Adsorbent Dosage

Adsorbent masses of 0.005, 0.01, 0.02, and 0.04 g Mg-Al LDH were added to separate 200 mL conical flasks containing 40 mg L⁻¹ HEDP solution (pH 7). The mixtures were agitated in a temperature-controlled shaker at 25 °C and 150 rpm for 3 h. The remaining HEDP concentration was quantified through the ammonium molybdate spectrophotometric approach using a UV-Vis spectrophotometer (UVmini-1285, Shimadzu, Kyoto, Japan).

2.4.2. Influence of Initial HEDP ConcentrationI

HEDP solutions (5, 10, 20, 40, 80, and 100 mg L⁻¹) were adjusted to pH 7 and treated with 0.02 g Mg-Al LDH. The suspensions were stirred at 25 °C and 150 rpm for 3 h, filtered (0.22 μm membrane), and analyzed for residual HEDP.

2.4.3. Influence of Solution pH

A 40 mg L⁻¹ HEDP solution was adjusted to pH 3–11 and contacted with 0.02 g Mg-Al LDH under the same stirring conditions. The filtrates were analyzed for residual HEDP.

2.4.4. Influence of Coexisting Ions

To assess the effect of coexisting ions, 0.02 g of Mg-Al LDH was placed in 200 mL conical flasks along with 100 mL of 40 mg L⁻¹ HEDP solution (pH 7, adjusted with 0.1 M NaOH or HNO₃). Selected ions (NO₃⁻, Cl⁻, SO₄²⁻, CO₃²⁻, Ca²⁺) were introduced separately at 10 mg L⁻¹, with a control sample containing no additional ions. All samples were shaken at 25 °C and 150 rpm for 3 h. Post adsorption, each solution was filtered through a 0.22 μm membrane, and the residual HEDP was measured to compute the equilibrium adsorption capacity.

2.5. Adsorption Kinetic Model

2.5.1. Kinetics Experiment

A 0.02 g sample of Mg-Al LDH was placed in a 200 mL conical flask. A 100 mL volume of HEDP solution (40 mg L⁻¹, pH adjusted to 7 with 0.1 M NaOH/HNO₃) was introduced into the flask. The mixture was shaken at 150 rpm and 25 °C in a thermostatic water bath for 3 h. The residual HEDP concentration was monitored over time to determine the adsorption capacity and evaluate the kinetics of the process. The obtained data were modeled using the pseudo-first-order (Equation (2)), pseudo-second-order (Equation (3)), and Elovich (Equation (4)) kinetic equations:

$$\ln \left(q_e-q_t\right)=ln{q}_e- k_{1}t$$

(2)

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

(3)

$$q_t={\displaystyle \frac{1}{\beta }}[\ln \left(\alpha \beta \right)+\ln t]$$

(4)

where t represents adsorption time (min), $q_e$ and $q_t$ refer to the adsorption capacities (mg g⁻¹) at equilibrium and time t, respectively, $k_1$ (min⁻¹) and $k_2$ (g (mg·min)⁻¹) are the rate constants of the pseudo-first-order and pseudo-second-order models, α (g (g·min)⁻¹) denotes the initial adsorption rate, and β (mg g⁻¹) is the desorption constant.

2.5.2. Adsorption Isotherm Experiment

Mg-Al LDH (0.02 g) was weighed into a series of 200 mL conical flasks. HEDP solutions with concentrations between 5 and 100 mg L⁻¹ were prepared and adjusted to pH 7 using 0.1 M NaOH or HNO₃. Then, 100 mL of each solution was added to the flasks. The mixtures were agitated at 25 °C and 150 rpm for 3 h. After adsorption, the residual HEDP concentration was analyzed to compute the adsorption capacity under different initial concentrations. The equilibrium data were fitted with the Langmuir (Equation (5)) and Freundlich (Equation (6)) isotherm models:

$$q_e={\displaystyle \frac{q_m{K}_L{C}_e}{1+K_L{C}_e}}$$

(5)

$$q_e=K_F{C}_e^{1/n}$$

(6)

where $q_e$ is the equilibrium adsorption capacity of the adsorbent (mg g⁻¹), $C_e$ is the equilibrium concentration of the solution (mg L⁻¹), $K_L$ is the Langmuir adsorption constant (L mg⁻¹), $q_m$ is the maximum adsorption capacity of the adsorbent (mg g⁻¹), $K_F$ is the Freundlich isotherm adsorption constant (mg/g), n is the adsorption constant.

2.5.3. Adsorption Thermodynamics Experiment

The adsorption of HEDP (40 mg L⁻¹, pH 7) onto 0.02 g of Mg-Al LDH was conducted at 25, 35, and 45 °C. The experimental data were used to calculate thermodynamic parameters including Gibbs free energy change (ΔG°), enthalpy change (ΔH°), and entropy change (ΔS°) via the following equations:

$$\ln K_d={\displaystyle \frac{∆S°}{R}}-{\displaystyle \frac{∆H°}{RT}}$$

(7)

$$∆G°=-RTln{K}_d$$

(8)

$$K_c={\displaystyle \frac{C_0-C_e}{C_e}}\times {\displaystyle \frac{V}{m}}$$

(9)

$$K_d=1000K_c$$

(10)

$$∆G°=∆H°-T∆S°$$

(11)

where K_d and K_c are the thermodynamic constant and equilibrium constant, respectively. R is a constant (8.314 J (mol·K)⁻¹), T is the absolute temperature, in degrees Celsius, ΔS° is the standard adsorption entropy change, ΔG° is the standard adsorption free energy, and ΔH° is the standard adsorption enthalpy change.

3. Results and Discussion

3.1. Characterization of Mg-Al LDH

The structural properties of the as-synthesized Mg-Al LDH were characterized by XRD, SEM, FTIR, and BET techniques. The XRD pattern (Figure 1a) showed sharp and characteristic diffraction peaks at 11.5°, 23.0°, 34.5°, and 60.3°, which are indexed to the (003), (006), (009), and (110) crystal planes of Mg-Al LDH, respectively, confirming the successful formation of a layered double hydroxide structure [19,20,21]. SEM imaging (Figure 1b) revealed a characteristic lamellar morphology with wrinkled textures, consistent with typical Mg-Al LDH materials and further verifying successful synthesis [12]. The FTIR spectrum (Figure S1) displayed a broad absorption band around 3450 cm⁻¹, associated with -OH stretching vibrations from interlayer water molecules [22]. A peak near 1350 cm⁻¹ is related to NO₃⁻ stretching vibrations, and the band at 550 cm⁻¹ corresponds to metal-oxygen (M-O) vibrations [23,24]. The specific surface area, a key parameter influencing adsorption performance, was evaluated using N₂ adsorption–desorption measurements at −195.8 °C (Figure 1d) [25]. The isotherm was classified as type IV with an H3 hysteresis loop according to IUPAC, indicating a mesoporous structure. As summarized in

Table 1. BET determination results of Mg-Al LDH.

Adsorbent	Specific Surface Area (m2/g)	Pore Volume (cm3/g)	Pore Diameter (nm)
Mg-Al LDH	78.38	0.35	13.11

, BET analysis determined an average pore width of 13.11 nm, a specific surface area of 78.38 m² g⁻¹, and a pore volume of 0.35 cm³ g⁻¹, corroborating the mesoporous nature of the material. These results confirm the successful preparation of nitrate-intercalated Mg-Al LDH with high surface area.

3.2. HEDP Adsorption

3.2.1. Adsorption Kinetics

The adsorption kinetics of Mg-Al LDH for HEDP removal were investigated to assess its performance, with the results presented in Figure 2a. Using an initial HEDP concentration of 40 mg L⁻¹ at 293 K, the adsorbent significantly reduced the HEDP concentration by 74.0% within the first 5 min, achieving an adsorption capacity of 172.5 mg g⁻¹. This rapid uptake is likely due to the abundance of accessible active sites on the fresh Mg-Al LDH surface. After 30 min, the process reached equilibrium with a maximum adsorption capacity of 199.6 mg g⁻¹, as the available sites became saturated. The kinetic data were further analyzed using pseudo-first-order, pseudo-second-order, and Elovich models. As illustrated in Figure 2a and summarized in

Table 2. Adsorption kinetic fitting parameters of HEDP on Mg-Al LDH.

Mg-Al LDH	Pseudo-First-Order Kinetics			Pseudo-Second-Order Kinetics			Elovich
qe (exp)/ (mg·g−1)	qe (cal)/ (mg·g−1)	k1/ (min−1)	R2	qe (cal)/ (mg·g−1)	k2/ g·(mg·min)−1	R2	α	β	R2
199.58	190.32	1.933	0.959	194.92	0.0194	0.98	2.18	0.12	0.997

, the Elovich model yielded the highest correlation coefficient (R² = 0.99) at the given concentration, suggesting it best describes the adsorption behavior and supporting a chemisorption-dominated mechanism [26]. Typically, multiple mechanisms influence diffusion in adsorption systems.

To better understand diffusion behavior and rate-limiting steps, the intra-particle diffusion model was applied. As shown in Figure 2b, the entire adsorption process did not pass through the origin and can be divided into two stages. The fitted lines for both stages also failed to intersect the origin, indicating that intra-particle diffusion was not the sole rate-controlling step. These two stages suggest that the adsorption of HEDP onto Mg-Al LDH proceeds via: (i) a rapid adsorption stage, during which HEDP quickly occupies the available active sites upon initial contact, predominantly governed by liquid film diffusion, and (ii) an intra-particle diffusion equilibrium stage, where the adsorbate gradual diffusion into adsorbent pores and eventual equilibrium, characterized by a slower adsorption rate.

3.2.2. Adsorption Isotherm

Adsorption isotherms provide a useful method for studying how adsorbates distribute between the adsorbent and the liquid phase once equilibrium is reached. The high HEDP adsorption performance of the Mg-Al LDH sample was analyzed by fitting the data to Langmuir and Freundlich isotherm models, with the resulting parameters illustrated in Figure 3 and summarized in

Table 3. Isothermal adsorption model fitting results of Mg-Al LDH for HEDP adsorption.

Models	Parameters	298 K	308 K	318 K
Langumir	qm (mg·g−1)	279.00	174.14	130.15
	KL (L·mg−1)	0.995	1.096	3.131
	R2	0.948	0.903	0.967
Freundlich	1/n	0.252	0.168	0.0724
	KF (mg g−1)	109.854	91.572	98.177
	R2	0.824	0.923	0.696
Redlich-Peterson	n	1.216	0.887	0.959
	KRp	198.396	886.177	799.656
	R2	0.947	0.908	0.914

. Comparative evaluation across different temperatures indicated that the Langmuir model yielded higher correlation coefficients (R² = 0.903–0.967) than the Freundlich and Redlich-Peterson models [27], suggesting that it more accurately represents the adsorption behavior of HEDP on Mg-Al LDH. This supports the view that adsorption occurs mainly through monolayer coverage on homogeneous surfaces, consistent with chemisorption as the dominant mechanism. In addition, the Langmuir constant K fell within a relatively low range (0.995–3.131) compared to values from the Freundlich (91.572–109.854) and Redlich-Peterson models (198.396–886.177), further affirming the suitability of the Langmuir model for this system [28]. The maximum adsorption capacity derived from the Langmuir isotherm represents the theoretical saturation uptake at each temperature. It was observed that the adsorption capacity of Mg-Al LDH for HEDP declined as temperature increased, from 279.0 mg g⁻¹ at 298 K to 130.1 mg g⁻¹ at 318 K, indicating that the process is more favorable at lower temperatures.

3.2.3. Adsorption Thermodynamics

Thermodynamic analysis provides valuable information regarding the energy variations and fundamental mechanisms involved in the adsorption process. Using the experimental data from Section 3.2.2 and Equations (8)–(10), the equilibrium distribution coefficient (K_d) was computed and linearly regressed against 1000/T. From the slope and intercept of the fitted line, the values of ΔH° and ΔS° were derived, which were subsequently applied in Equation (11) to determine the Gibbs free energy change (ΔG°). The fitting curve is displayed in Figure 3d, and all resulting thermodynamic parameters are summarized in

Table 4. Thermodynamic parameters for the adsorption of HEDP onto Mg-Al LDH.

T (K)	∆G° (kJ mol−1)	∆S° (J (mol·K)−1)	∆H° (kJ mol−1)
298	−11.98	−345.11	−114.82
308	−8.53
318	−5.08

.

As indicated in Table 4, the values of ΔG° were negative across all temperatures studied and became more negative as temperature decreased, suggesting that the adsorption of HEDP onto Mg-Al LDH is a spontaneous process. The reaction is less favorable at higher temperatures, with improved adsorption observed under lower temperature conditions. The negative ΔH° value confirms that the adsorption is exothermic, aligning with the observed inhibitory effect of temperature increase. Furthermore, the negative ΔS° implies a reduction in randomness at the solid-solution interface during adsorption, indicating an entropy-decreasing process.

3.3. Effect of Operating Parameters on HEDP Removal by Mg-Al LDH

The maximum adsorption capacity of Mg-Al LDH for HEDP was evaluated at 298 K across a range of initial HEDP concentrations (5–100 mg L⁻¹), as depicted in Figure 4a. With increasing HEDP concentration from 5 to 100 mg L⁻¹, the removal efficiency decreased from 100% to 51%, which is likely due to the limited number of active sites available at higher solute concentrations. At concentrations below 40 mg L⁻¹, the removal efficiency remained high (>96%), while a further increase beyond 40 mg L⁻¹ resulted in a notable reduction to approximately 50%, indicating site saturation. In contrast, the adsorption capacity rose from 23.8 to 276.0 mg g⁻¹, demonstrating a strong affinity between Mg-Al LDH and HEDP.

As illustrated in Figure 4b, when the adsorbent dosage was below 0.2 g L⁻¹, the adsorption capacity increased from 158.8 to 199.6 mg g⁻¹, accompanied by an improvement in removal efficiency from 20% to 97%, owing to the greater number of available active sites. However, further increasing the dosage to 0.4 g L⁻¹ led to a noticeable decline in adsorption capacity to 110.8 mg g⁻¹, likely due to particle aggregation reducing the accessible surface area. Despite this, removal efficiency remained high (~98%), as ample sites were still present to adsorb HEDP.

Solution pH significantly influences both the surface charge of Mg-Al LDH and the ionization state of HEDP, thereby affecting adsorption behavior. As shown in Figure 4c, adsorption capacity initially increased and then decreased with rising pH. When pH increased from 3.0 to 7.0, capacity increased from 153.52 to 199.58 mg g⁻¹. At low pH (<5.0), partial dissolution of the LDH may reduce active sites [29]. Near neutral pH, HEDP is primarily present as higher-charged anions (Figure 5a), facilitating adsorption through electrostatic attraction (pH_pzc of Mg-Al LDH = 11.6, Figure 5b). Beyond this pH, the adsorption capacity decreases gradually, which can be attributed to the increasing concentration of OH⁻ ions in solution competing with HEDP for adsorption sites.

Moreover, natural water bodies and wastewater often contain abundant anions and cations (e.g., CO₃²⁻, NO₃⁻, Cl⁻, Ca²⁺), which may interfere with the adsorption and removal of HEDP by Mg-Al LDH. Therefore, the effects of these coexisting ions on HEDP removal by Mg-Al LDH were investigated, and the results are shown in Figure 4d. As observed, anions exhibited a much stronger inhibitory effect on HEDP removal than cations, with the order of interference being CO₃²⁻ > SO₄²⁻ > Cl⁻ > NO₃⁻ > Ca²⁺. In the presence of Ca²⁺, the HEDP removal efficiency of Mg-Al LDH decreased to 85%, likely due to the formation of soluble Ca-HEDP complexes that remain in the aqueous phase. In contrast, inorganic anions carry negative charges and can compete with HEDP for adsorption sites, leading to reduced removal efficiency [30]. Among them, CO₃²⁻ showed the strongest inhibitory effect, with the final adsorption capacity dropping to 124.8 mg g⁻¹ and the removal efficiency to 60.1%. This pronounced effect can be attributed to CO₃²⁻ readily forming mineral precipitates with surface Mg on LDH, thereby eliminating active adsorption sites [31]. Furthermore, CO₃²⁻ in solution can hydrolyze to produce OH⁻, which further competes with HEDP for adsorption [10]. In the future, two approaches may be employed to mitigate the influence of CO₃²⁻: (i) applying simple pretreatment steps, such as acidification or aeration–degassing, to reduce the carbonate concentration in wastewater; and (ii) modifying the LDH material, for example, through surface functionalization or by incorporating transition metals and organic ligands to enhance its selectivity toward phosphonate groups over carbonate.

3.4. Adsorption Mechanism

The adsorption kinetics and isotherm analysis indicated that HEDP removal was predominantly governed by chemisorption, along with liquid film within a monolayer adsorption process. In Section 3.3, we found that electrostatic attraction between Mg-Al LDH and HEDP facilitated HEDP removal, but it was not the sole driving force for adsorption. To further investigate other mechanisms involved in HEDP adsorption on Mg-Al LDH, the surface properties of Mg-Al LDH before and after adsorption were examined using XRD and FTIR analyses. Figure 1a presents the XRD patterns of Mg-Al LDH before and after adsorption. No new mineral phases were detected after adsorption, indicating that HEDP was not removed via surface precipitation. According to Bragg’s law, the interlayer spacing remained constant at 0.802 nm, indicating that ion exchange did not occur between HEDP and interlayer anions. These results suggest that adsorption took place primarily on the surface. Therefore, in this study, the adsorption process is likely dominated by surface adsorption. As shown in the FTIR spectra (Figure 1c), a decrease in intensity of the -OH stretching band at 3500 cm⁻¹ after HEDP adsorption, implying possible complexation via hydroxyl ligand exchange. The nitrate vibration band showed negligible change, supporting that interlayer NO₃⁻ ions were not exchanged. New peaks emerged at 1050 cm⁻¹ and 590 cm⁻¹, corresponding to P-O asymmetric stretching and O-P-O bending vibrations from HEDP, respectively [32,33]. A shift in the M-O vibration peak near 544 cm⁻¹ was also observed, indicating the involvement of metal atoms in the Mg-Al LDH layers during adsorption. Given the high affinity of HEDP toward Mg and Al ions (log K = 22.6 and log K = 22.7, respectively), the adsorption is largely driven by surface complexation at metal sites, supplemented by electrostatic attraction [34,35]. A proposed mechanism for HEDP removal by Mg-Al LDH is illustrated in Figure 6.

3.5. Comparison with Other Adsorbents and Reusability

To better evaluate the adsorption performance of Mg-Al LDH, its maximum adsorption capacity for HEDP was compared with those of other reported adsorbents (Table S1) [10,36,37,38,39,40]. As shown, Mg-Al LDH exhibits a much higher adsorption capacity of 279 mg g⁻¹ (82.8 mg-P g⁻¹) than most of the previously reported adsorbents, such as Nd-based nanocomposites (12.2 mg-P g⁻¹), Eu-MOF (37.31 mg-P g⁻¹), kaolin/lanthanum carbonate composites (4.93 mg-P g⁻¹), and zirconium-loaded zeolite (3.77 mg-P g⁻¹). These results clearly confirm that Mg-Al LDH possesses superior affinity toward HEDP, highlighting its promise as a highly efficient material for phosphonate removal.

In addition to its superior capacity, the reusability of LDH was also evaluated, as it represents a key factor for practical application. After five consecutive adsorption–desorption cycles, the adsorption capacity decreased only slightly from 199.6 mg/g to 172.3 mg/g (Figure 7), indicating that LDH maintains good stability and recyclability. These results suggest that LDH is not only highly efficient but also cost-effective and sustainable for repeated use in large-scale applications.

4. Conclusions

This study demonstrated that Mg-Al LDH is a highly effective adsorbent for the removal of the HEDP. Under optimized conditions, the adsorption efficiency exceeded 95%, with a maximum adsorption capacity of 276.0 mg g⁻¹ at 298 K. Kinetic and isotherm analyses revealed that the adsorption process followed Elovich kinetic model and Freundlich isotherm model, indicating monolayer chemisorption dominated by electrostatic interactions. Characterization techniques, including XRD, FTIR, and zeta potential measurements, confirmed that HEDP adsorption occurred primarily through surface complexation with metal sites (Mg and Al) rather than interlayer ion exchange. The adsorption performance was influenced by solution pH, coexisting ions, and temperature, with optimal removal observed at neutral pH and lower temperatures. Competitive anions, particularly CO₃²⁻, significantly inhibited HEDP adsorption due to site competition and mineral precipitation. These findings underscore the promise of Mg-Al LDH as a cost-effective and sustainable material for mitigating phosphonate pollution in wastewater, offering a viable solution for industries reliant on HEDP while addressing environmental concerns. Future research could explore large-scale applications and regeneration techniques to further enhance practicality.