Enhancing the Immobilization of Hexavalent Chromium by the Interlayer Anion Adsorption of the Brucite-Transformed LDH in the Presence of Aluminum Ions

Abstract

Current studies of chromium adsorption kinetics at the solid–liquid interface often neglect the influence of coexisting ions in complex wastewaters. Thus, it is critical to explore the hexavalent chromium Cr(VI) adsorption kinetics of solid-phase brucite (Mg(OH)₂) in liquid-phase wastewater containing coexisting aluminum ions (Al(III)). This paper reveals that the presence of Al(III) significantly enhanced the Cr(VI) adsorption efficiency onto Mg(OH)₂, with a peak of up to 91% compared to 5% for the absence of Al(III). The main reason for this enhancement was the initial surface ternary complexation of Mg(OH)₂ and the cationic (Al(III)) isomorphic substitution to form Mg(II)-Al(III) layered double hydroxides (LDH), which also indicates a solid-phase transition on the surface of Mg(OH)₂, which led to electrostatic adsorption in the gallery and made Cr(VI) immobilized and not readily released. Further calculation and analysis of the adsorption energy confirmed the mechanism of Cr(VI) adsorption. It was also concluded that Cr(VI) migration in Mg(OH)₂-containing minerals was affected by the phase transformation of solids in the presence of Al(III). Hence, this study not only reveals the adsorption mechanism during the treatment of composite pollutant wastewater but also provides the methodological reference for brucite synergistic adsorption to remove heavy metal ions and purify and treat complex polluted wastewater.

1. Introduction

Chromite [(Fe, Mg)Cr₂O₄], as one of the main chromium-containing minerals in the lithosphere, is popularly mined for the production of sodium dichromate (Na₂Cr₂O₇) and other compounds containing hexavalent Cr(VI) in many chemical industries, including pigment, textile, electroplating, leather, etc. The processes of chromite smelting, industrial manufacture, and combustion could cause a great deal of emission of Cr(VI)-containing exhaust gas, wastewater, and waste residues. As a consequence of the cycling of Earth materials, heavy metal ions Cr(VI) could migrate and transform in the hydrosphere, pedosphere, and lithosphere [1,2]. Due to the fact that Cr(VI) could pose great threats to people’s health, resulting in skin allergies and cancer [3], and that the easy migration property of Cr(VI) could lead to sustained harm to the natural environment, studies concerning the removal and migration of Cr(VI) in the environment are of great significance and have attracted continued attention [4,5,6].

Since chromite is composed primarily of spinel FeCr₂O₄, research in this regard is mainly Fe-based, focusing on such hydroxides as Goethite, Hematite [7], and Fe-hydrotalcite [8]. On the basis of the adsorption of Fe-based hydroxide, it has already been discovered to be involved in the migration of Cr(VI) [9,10]. However, these studies neglected the effect of Mg and Al elements, which are also present in chromium-containing slag, sludge, and wastewater, on the adsorption migration of Cr(VI), such as Al₂O₃ and MgO, which are the main components of chromium-containing slag [11,12]. In addition, hydrogarnet, magnesium aluminum hydrotalcite (Mg-Al-HTs), and brucite (Mg(OH)₂) can be found in chromite ore processing residues (COPR), which are the main binding phases of Cr(VI), and their storage may lead to the dissolution of Cr(VI) with an environmental change [13,14], which in turn causes environmental pollution. Based on the above conclusions, it is crucial to study the effect of the presence of Al(III) on the adsorption of Cr(VI) by Mg(OH)₂. M. A. Gomez et al. also proposed that uranium is a heavy metal element similar to Cr where Al, Mg, or Mg-Al sodium hydroxide mineral phases are present [15]. These magnesium- and aluminum-containing minerals are important for the grinding process and associated sediment formation processes and can be effective in removing potential environmental contaminants from the sedimentary system [15]. Specifically, solid-phase brucite (Mg(OH)₂) is a natural mineral with adsorption properties [16,17], and its adsorption and migration in aqueous geochemical systems and human activities are often accompanied by other metal ions [18,19].

Also, brucite (Mg(OH)₂) can be used as a green adsorbent in the wastewater treatment process. Meanwhile, during the adsorption of brucite (Mg(OH)₂), its solid-phase interface will react with the liquid-phase interface by interfacial interaction and transform the solid-phase brucite structure, thus affecting the capacity and efficiency of Cr(VI) adsorption. Hongbo Lu et al. proposed that strong interfacial interaction will occur during the adsorption and desorption of heavy metal ions Sb(III) and solid-phase natural mineral birnessite, and the solid phase transformation will occur, which is the transformation of birnessite to vernadite [20]. It is also proposed in another paper that anion exchange occurs during the co-adsorption of Cr(VI) and Cu(II) by Cl-LDH, which is the isomeric substitution of Mg(II) with Cu(II), and a solid-phase transformation occurs to form Cu₂Cl(OH)₃, which in turn improves the adsorption of Cr(VI) [21]. Compared to the previously mentioned substances, brucite actually undergoes more phase transitions in the interlayer anion exchange, which is the replacement of Mg(II) in Mg(OH)₂ by Al(III), which in turn forms LDH and affects the adsorption of Cr(VI).

The performance of nano-Mg(OH)₂ for the adsorption of metal oxyanions such as U(VI) and Cr(VI) has been studied [17,22]. One of the studies, through molecular dynamics simulation, showed that the adsorption of U(VI) by Mg(OH)₂ is mainly a function of the surface hydroxyl group [22]. In the treatment of low concentrations of Cr(VI)-containing wastewater, the study of Liu et al. used Mg(OH)₂ for Cr(VI) adsorption and reached an equilibrium adsorption capacity (1.55 mg/g) after 3 h of adsorption [17], but the coexistence of Al(III) in the wastewater should also be taken into account, the importance of which has been mentioned in this paper. Meanwhile, the process of Cr(VI) adsorption by Mg(OH)₂ in the presence of Al(III) ions has rarely received attention in the current studies. This is especially true considering that the kinetics of chromium adsorption at the solid–liquid interface tend to neglect the influence of coexisting ions in the composite wastewater.

This research aimed to explore the adsorption process and kinetics of the synthesized Mg(OH)₂ nanosheets on Cr(VI)-containing wastewater in the presence of Al(III). It was found that a phase transformation existed in this adsorption process, whose mechanism should be explored. In addition, the less water-soluble brucite (Mg(OH)₂) could gradually dissolve and release OH⁻ [23,24], which means the pH value changes over time in the natural environment. Therefore, considering the factors affecting the adsorption and migration of Cr(VI) onto Mg(OH)₂ adsorbent, the pH variation was monitored, and the migration of Cr(VI) at different pH values was recorded and analyzed. Overall, research on the adsorption and phase transition mechanisms on Mg(OH)₂-Cr(VI)-containing solution surfaces, which can also be regarded as the adsorption process of heavy metal ions (Cr(VI)) by the natural mineral brucite with Mg(OH)₂ as the main component, can help to elucidate the study of the environmental pollutants Cr(VI) in eco-chemical circulation and can even be used in actual wastewater purification treatment [17].

2. Materials and Methods

2.1. Materials and Reagents

All chemicals were of analytical reagent quality and were not purified any further. Chemical Reagent Beijing Co., Ltd. (Beijing, China) supplied KCl (99.5%, AR, CAS No. 7447-40-7), Na₂SO₄ (99%, AR, CAS No. 7757-82-6), Na₂CO₃ (99.5%, AR, CAS No. 497-19-8), NaHCO₃ (99.5%, AR, CAS No. 144-55-8), and NaOH (95%, AR, CAS No. 1310-73-2). And magnesium chloride hexahydrate (MgCl₂•6H₂O (95%, AR, CAS No. 7791-18-6)), aluminum chloride hexahydrate (AlCl₃•6H₂O (98%, AR, CAS No. 7784-13-6)) and non-aqueous aluminum nitrate (Al(NO₃)₃•9H₂O (95%, AR, CAS No. 7784-27-2)) were utilized as the sources of metal cations. Additionally, throughout the examinations, ultrapure water with a resistivity of 18 MΩcm was used.

2.2. Synthesis of Mg(OH)₂ Nano-Sheets

To synthesize Mg(OH)₂ nano-sheets, 20.5 g of MgCl₂•6H₂O were dissolved in a solution of 25 mL of anhydrous alcohol (99.5%, AR, CAS No. 64-17-5) and 125 mL of ultrapure water; 6.30 g of NaOH were dissolved in 100 mL of ultrapure water; and then the NaOH solution was steadily introduced and stirred for 10 min in the MgCl₂ solution. Then, the precipitated phase developed after the suspension had matured at 65 °C for 6 h. To eliminate excessive NaOH, the product was centrifuged at 10,000 rpm for 10 min, rinsed three times with ultrapure water, and the precipitate was collected and dried. Additionally, the Brunauer–Emmett–Teller (BET) examinations revealed that Mg(OH)₂ had a specific surface area of 42.6 m²/g (Figure S1).

2.3. Cr(VI) Adsorption Experiments

The migration process of Cr(VI) in magnesium aluminum hydrotalcite was investigated by adsorption experiments. In a kinetics study, a mixed 150 mL solution of Cr(VI) (5 mg/L) and Al(III) ions (1 mM) was generated by dissolving potassium dichromate (K₂Cr₂O₇, 99.5%, AR, CAS No. 7778-50-9) and aluminum chloride hexahydrate (AlCl₃•6H₂O) into ultrapure water. The above obtained 0.1 g of Mg(OH)₂ substances were added to the mixed solution and then shaken well on a thermostatic shaker (SYC-B). After a specified contact time for adsorption (30 min and 24 h), the mixture was centrifuged at 5000 rpm for 5 min, and the supernatant was then filtered through a filter membrane with a particle size of 0.22-μm. The elemental content of the filtrate was analyzed. A tiny portion of these samples was air-dried before being analyzed with X-ray diffraction (XRD) and Fourier transform infrared spectroscopy (FT-IR). Two parallel experiments were carried out. The removal rate R (%) of Cr(VI) was represented by Equation (1), and the equilibrium adsorption amount q_e (mg/g) of Cr(VI) was represented by Equation (2) [25].

$$R\left(\%\right)=\frac{\left(C_0-C_t\right)}{C_0}\times 100$$

(1)

$$q_e=\frac{\left(C_0-C_e\right)V}{m}$$

(2)

where the original concentration of adsorbate in the solution is expressed as C₀ (mg/L). The C_t (mg/L) represents the concentration of Cr(VI) in the adsorbent solution at a fixed time t, and the C_e (mg/L) represents the equilibrium concentration of Cr(VI) at adsorption equilibrium. V (mL) denotes the amount of the solution, and m (g) denotes the mass of the adsorbent.

2.4. Regeneration Experiment of Cr(VI)

Three anionic desorbents (NaHCO₃, Na₂CO₃, and Na₂SO₄) were used to desorb the solid samples obtained after adsorption experiments for each of them (Table S1), and NaHCO₃ was selected for the regeneration experiments.

The specific brief steps were as follows: Firstly, the Mg(OH)₂ solid sample containing Cr(VI) obtained by centrifugal drying after the adsorption reaction was collected. Next, 0.1 g of solid sample was added to a 150 mL solution of NaHCO₃ (1 mM), and the sample was desorbed by shaking on a thermostatic shaker (SYC-B) for 12 h. After centrifugation at 5000 rpm for 5 min, the precipitate obtained was dried at room temperature and then preserved. The solid product was preserved for the regeneration treatment of Mg(OH)₂, and the specific experimental steps were as in 2.3 for the adsorption of Cr(VI) by Mg(OH)₂ in the presence of Al(III).

2.5. Analytical Methods

The BET analysis method (ASAP2020 HD88, Micromeritics, Norcross, GA, USA) was used to measure the specific surface area of the Mg(OH)₂ sample. XRD analysis of samples before and after adsorption was performed on a computer-controlled X-ray diffractometer (X’ Pert³ Powder, PANalytical, Almelo, The Netherlands) to determine the formation of layered Mg(II)-Al(III)-hydroxides. Continuous screening with Cu-Kα radiation (40 kV, 40 mA) was used to capture the diffraction data.

FT-IR (Perkin–Elmer, Waltham, MA, USA) and the KBr particle method were used to identify any chemical bonding and surface functional group alterations in the samples. The spectra range was 4000–400 cm⁻¹ with a resolution of 4 cm⁻¹. The samples were examined using a transmission electron microscope (TEM, HT7700, Hitachi, Tokyo, Japan) to unravel their morphology, and the surface elemental composition of the samples was analyzed using its equipped Energy Dispersive X-ray Spectroscopy (EDX).

Furthermore, the concentration of Cr(VI) was determined by the spectrophotometric method (GB 7467 and GB/T 15555.4-1995, China EPA) [26,27]. Thus, Cr(VI) was treated in an acidic medium with 1,5-diphenylcarbohydrazide, and the resulting product was measured using a UV-vis spectrophotometer (TU-1810, Persee, Beijing, China) at 540 nm. The concentrations of Mg(II) and Al(III) in the filtrate were determined just using an inductively coupled plasma optical emission spectrometer (OES, OPTIMA 8300, Perkin–Elmer).

3. Results and Discussion

3.1. The Adsorption of Cr(VI) in Mg(OH)₂

Adsorption kinetic experiments can reflect the relationship between the change in Cr(VI) removal rate with increasing reaction time and are used to simulate the migration of Cr(VI) in brucite (Mg(OH)₂). In this experiment, Mg(OH)₂ nano-sheets were prepared to adsorb Cr(VI) in the absence and presence (0.5 mM or 1 mM) of Al(III), respectively. As illustrated in Figure 1a, the percentage removal of Cr(VI) by Mg(OH)₂ was only 5% in the absence of Al(III). According to the BET analysis, the specific surface area of Mg(OH)₂ was only 42.6 m²/g (Figure S1), and the lower specific surface area was considered partly responsible for the lower removal rate without the addition of Al(III). However, in the presence of Al(III) (0.5 mM and 1 mM), the specific surface area of Mg(OH)₂ was 72.5 m²/g and 69.6 m²/g (Figure S2), respectively, considering that the structure of the solid sample changed in the presence of Al(III), which affected the specific surface area of Mg(OH)₂. Meanwhile, the addition of Al(III) (0.5 mM or 1 mM) had a discernible effect on the kinetics and extent of Cr(VI) adsorption, with a marked increase in the removal of Cr(VI). In the presence of 1 mM Al(III), the removal rate of Cr(VI) reached more than 90% (91%) in 40 min, which was significantly greater compared to the absence of Al(III) (only 5%). Meanwhile, according to the comparison of the adsorption amounts in different references, it was found that the adsorption amount of Mg(OH)₂ on Cr(VI) was only 1.6 mg/g in the study of Lin et al. and only 0.4 mg/g in the present study (Table S2) [17]. In the presence of Al(III), the amount of Cr(VI) adsorption by Mg(OH)₂ solid samples was increased to 3 mg/g in this study, indicating that Al(III) has a critical influence on Cr(VI) adsorption by Mg(OH)₂. Therefore, the migration reduction of Cr(VI) in the presence of brucite is considered to be due to the effect of the addition of Al(III).

However, further analysis of the kinetics of Cr(VI) adsorption in the presence of Al(III) revealed an interesting phenomenon. As shown in Figure 1a,b, the Cr(VI) adsorption rate peaked at 91% after about 40 min when the initial concentration of Al(III) was 1 mM and then steadily decreased to an equilibrium level of around 40%. Specifically, the whole process was divided into two stages, including adsorption (stage 1) and desorption (stage 2). Meanwhile, the powdered solid samples after adsorption were collected by centrifugation under different reaction conditions in order to observe the color of the samples after reactions. As shown in Figure 1c, Cr(VI) adsorbed onto the surface of Mg(OH)₂, which caused the originally white Mg(OH)₂ samples (Figure 1c(①)) to turn yellow. The sample reacted in the presence of Al(III), and the color was darker at the maximum adsorption amount (Figure 1c(②)) compared with the equilibrium adsorption condition (Figure 1c(③)), and the color was darker in the presence of Al(III) than without the addition of Al (III) (Figure 1c(④)). The overall appearance of the samples demonstrated that the variation of Cr(VI) adsorption under different reaction conditions was consistent with the experimental data shown in Figure 1a. To summarize, the addition of Al(III) to the reaction process improved the adsorption efficiency of Cr(VI), accompanied by the desorption of Cr(VI) under specific circumstances. This indicated that the migration of Cr(VI) in the environment is significantly influenced by Al(III).

Then, in order to further explore the behaviors of stages 1 and 2, the adsorption and desorption kinetics data were fitted to pseudo-first-order and pseudo-second-order models, as shown in Figure 2. Their linear equations are listed as follows [28,29]:

$$\ln \left(q_e-q_t\right)=\ln q_e-k_1$$

(3)

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

(4)

where q_t (mg/g) denotes the adsorption capacity at moment t and q_e (mg/g) denotes the adsorption capacity at equilibrium. k₁ (min⁻¹) is the pseudo-first-order model rate constant, and k₂ (g/(mg·min)) is the pseudo-second-order model rate constant. The effects of contact time on the nonlinear curves of the pseudo-second-order and pseudo-first-order kinetic models for adsorption (stage 1) and desorption (stage 2) onto Mg(OH)₂ are illustrated in Figure 2a,b, respectively. Since the closer the R² is to 1, the better the fit, the adsorption data (stage 1) best matched the pseudo-second-order kinetic model for Cr (VI) adsorption (R² = 0.995) based on R² values and a comparison of experimental and calculated q_e values [30]. Furthermore, the desorption data (stage 2) fit best with the first-order kinetic models for Cr(VI) desorption (R² = 0.929). It showed that there were chemical changes of bonding and non-homogeneous phase changes in the formation of brucite and the adsorption of Cr(VI), rather than just diffusion. However, the desorption data (stage 2) were best fitted with the first-order kinetic models of Cr(VI) desorption (R² = 0.929), so stage 2 was basically controlled by diffusion, and part of the desorbed Cr(VI) was adsorbed onto the Mg(OH)₂ surface by weak complexation.

3.2. The Formation of LDH in the Cr(VI) Adsorption

The structure of the powder samples was further characterized by centrifugation in chemical kinetic experiments to explain the aforementioned phenomenon during the reaction. All diffraction peaks in the X-ray diffraction (XRD) patterns of the prepared Mg(OH)₂ sample (Figure 3a, PDF card 01-084-2163) could well demonstrate that the sample is the magnesium hydroxide in Figure 3a. Furthermore, the XRD pattern of these samples already exhibited two diffraction peaks at 11.5° and 23.3° as the Cr(VI) adsorption reaction with Al(III) proceeded, compared to the samples without Al(III) (Figure 3b, Figures S3 and S4). This pattern was the same as that of the typical LDHs, that is, the (003) and (006) crystal facet peaks, indicating the formation of Mg(II)-Al(III)-layered double hydroxide (LDH) by Mg(OH)₂ precipitation in the presence of Al(III). Then, the TEM images revealed the structure of the prepared (Figure 3c(①)) and reaction adsorbed (Figure 3c(②)) Mg(OH)₂, both as hexagonal nano-sheets, and comparison of the TEM images also showed that the LDH samples formed on the surface of Mg(OH)₂ still remained a sheet-like structure, and the stacking between the layers seemed to be tighter. In addition, FT-IR analysis was an effective approach to the characterization of the prepared Mg(OH)₂ solid sample, the solid samples in different reaction times of Cr(VI) adsorption with Al(III), and Mg(II)-Al(III)-LDH (Figure 3d and Figure S5). As shown in Figure 3d, as the adsorption time increased, all of the characteristic adsorption bands of Mg(OH)₂ remained in the powder samples. In all spectra, a stronger, wider, and slightly enhanced adsorption band at 3450 cm⁻¹ was found for these samples after the adsorption of Cr(VI) with Al(III) present. This property is consistent with Mg-Al-LDHs. The energy band characteristic represented the hydroxyl stretching vibration peak, which was attributed to the hydroxyl stretching between the brucite-like layers and the water molecular layer [8]. Due to the formed LDH, strong hydrogen bonds were formed between the LDH layer and interlayer anions, the electrostatic attraction was enhanced, the O-H bonds were strengthened, and the energy of stretching vibration increased. Additionally, it has been reported that free chromate ions show a stretching pattern of Cr-O at 890 cm⁻¹ [8], and we could obtain the fact of chromate presence in the samples by measuring the difference in Cr(VI) concentration in the reaction solution before and after the adsorption reaction. However, the samples in this study did not show a significant characteristic band of chromate near 890 cm⁻¹ after the adsorption of Cr(VI) with Al(III), considering the low adsorption of Mg(OH)₂ samples after the adsorption equilibrium. Based on other related studies, the stretching vibration of Cr-O when Cr(VI) was adsorbed onto LDHNSs led to the appearance of a new adsorption band in the spectrum at 885 cm⁻¹ [31]. Compared to Mg(II), Al(III) has a higher charge and smaller ionic radius, and after Al(III) replaces Mg(OH)₂ on Mg(II), the positively charged solid-like adsorbed anionic chromate is formed (Figure 3b). Moreover, according to all the spectra, an adsorption band appeared at 1634 cm⁻¹, which can be ascribed to the δ(H-OH) vibration mode [32]. In addition, it was also found that the specific surface area of Mg(OH)₂ increased to some extent (42.6 m²/g to 69.6 m²/g) after the addition of Al(III) (1 mM) by the conclusion of the BET data (Figures S1 and S2). In all, the structural characterization of the samples revealed that the process of Cr(VI) adsorption by Mg(OH)₂ is accompanied by the formation of Mg-Al-LDHs in the presence of Al(III). Referring to the papers that have been reported, LDH prepared for Cr(VI) adsorption did have a higher equilibrium adsorption amount (q_e = 12.2 mg/g, Table S2), and in this study, the LDH formed on the surface of the Mg(OH)₂ solid sample in the presence of Al(III) promoted the adsorption of Cr(VI). When Cr(VI) is present in minerals containing Mg(OH)₂, LDH is formed whenever Al(III) is present, resulting in the migratory adsorption of Cr(VI).

Further monitoring of the content of Mg(II), Al(III), and pH in the solution during the reaction revealed that pH had obviously changed from 4 to neutral gradually (Figure 4a, [Mg(OH)₂] = 0.67 g/L, [Al(III)]_initial = 1 mM, [Cr(VI)]_initial = 5 mg/L) with the progress of the reaction. Xu et al. proposed that the release of Mg(OH)₂ increased the pH of the neutral solution [33]. Therefore, the presence of a small amount of Mg(OH)₂ dissolved during the experiment led to an increase in the pH of this reaction. According to previous research [14], COPR is enriched with brucite, an Mg(OH)₂ mineral, which may also be responsible for its high alkalinity in the natural environment. In addition, the initial concentration of 1 mM Al(III) in the solution was basically removed with a removal rate of 93%, and equilibrium was reached (Figure 4a). The Ksp of aluminum hydroxide (Al(OH)₃) calculated that the precipitate was formed at a concentration of 4.6 mM Al(III) when pH = 4. Therefore, at the beginning of the reaction (pH = 4), Al(OH)₃ precipitation did not occur instantly (Figure 4a), which was also consistent with the phenomenon that there was no characteristic peak of Al(OH)₃ in the XRD patterns (Figure 3a). Initially, Mg(OH)₂ in this reaction system had some hydroxyl groups that could act as functional groups to chemically combine with Al(III) to form inner sphere complexes [34]. Meanwhile, there might be two types of ternary surface complexes (Equations (5) and (6)). Additionally, as the reaction proceeded, it was also detected that the Mg(II) content gradually increased and reached equilibrium (about 37 ppm). The molar ratio of Mg(II) and removed Al(III) in the solution was about 3:2, and Mg(II) was released due to the dissolution of part of the Mg(OH)₂ surface by the substitution of Al(III). In addition, EDX also indicated that Al(III) and Cr(VI) were present on the surface of Mg(OH)₂ after 24 h of adsorption (Figure S6). As the Mg(II) was released and pH increased, Mg(II) was substituted with Al(III) adsorbed on the inner layer, resulting in a positively charged interlayer structure, which was further neutralized by the interlayer anions (Cr(VI), Cl⁻) (Equations (7) and (8)). Therefore, Cr(VI) was continuously removed from the solution and formed Mg(II)-Al(III)-LDH with HCrO₄⁻.

$$\left[\mathrm{Mg}\right(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}]\mathrm{Sur}\text{-}\mathrm{OH} + {\mathrm{HCrO}}_{4^{-}} +{\mathrm{Al}}^{3+} \rightleftarrows [\mathrm{Mg}\left(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}\right]\mathrm{Sur}\text{-}\mathrm{OAl}\text{-}{\mathrm{HCrO}}_{4^{-}} + {\mathrm{H}}^{+}$$

(5)

$$\left[\mathrm{Mg}\right(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}]\mathrm{Sur}\text{-}\mathrm{OH} + {\mathrm{HCrO}}_{4^{-}} +{\mathrm{Al}}^{3+} \rightleftarrows [\mathrm{Mg}\left(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}\right]\mathrm{Sur}\text{-}{\mathrm{HCrO}}_4\text{-}{\mathrm{Al}}^{3+} + \mathrm{O}{\mathrm{H}}^{-}$$

(6)

$$\left[\mathrm{Mg}\right(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}]\mathrm{Sur} + {\mathrm{Al}}^{3+} + \mathrm{O}{\mathrm{H}}^{-}+ {\mathrm{HCrO}}_{4^{-}} + {\mathrm{H}}_2\mathrm{O}\to \mathrm{LDH}[\mathrm{M}{\mathrm{g}}_{1-\mathrm{x}}\mathrm{A}{\mathrm{l}}_{\mathrm{x}}\left(\mathrm{OH}{)}_2\right({\mathrm{HCrO}}_{4^{-}}{)}_{\mathrm{x}} \mathrm{n}{\mathrm{H}}_2\mathrm{O}]$$

(7)

$$\left[\mathrm{Mg}\right(\mathrm{OH}{)}_{2\left(\mathrm{S}\right)}]\mathrm{Sur} + {\mathrm{Al}}^{3+}+ \mathrm{O}{\mathrm{H}}^{-}+ \mathrm{C}{\mathrm{l}}^{-} + {\mathrm{H}}_2\mathrm{O}\to \mathrm{LDH}[\mathrm{M}{\mathrm{g}}_{1-\mathrm{x}}\mathrm{A}{\mathrm{l}}_{\mathrm{x}}\left(\mathrm{OH}{)}_2\right({\mathrm{Cl}}^{-}{)}_{\mathrm{x}} \mathrm{n}{\mathrm{H}}_2\mathrm{O}]$$

(8)

The Cr(VI) between the LDH layers was more stabilized [35] and more difficult to desorb, with the Cr(VI) removal efficiency at equilibrium remaining at 40% in Figure 1a. According to Zhang et al. [31], the total surface area and reaction rate might be increased due to the LDH nanosheets (LDHNSs) during Cr(VI) adsorption by exfoliated LDHNSs. During the adsorption reaction, Cr(VI) weakly interacted with the LDHNS binding sites. Taking all this into account, in this experiment, the Mg(II) on Mg(OH)₂ nanosheets was isomorphically substituted by Al(III), and the end product was equivalent to the LDHNSs, which subsequently formed LDH (Equation (7)). Therefore, for the adsorption of Cr(VI), the interlayer effect of the formed LDH structure might be stronger than the ternary complex on the surface of Mg(OH)₂, and thus Cr(VI) was difficult to desorbed. Meanwhile, in this study, the Cr(VI) removal efficiency reached a maximum at pH = 6 (Figure S7), which is consistent with the optimal pH of 6.0 described in the literature for the removal of Cr(VI) using the LDHNSs [31].

Furthermore, we also performed adsorbent regeneration experiments to desorb Mg(OH)₂ containing Cr(VI), using three anions separately (Table S1). After comparing the desorption efficiency, NaHCO₃ was chosen as the desorbent, and the desorbed solid samples were subjected to adsorption experiments again under the same conditions as those for the adsorption of Cr(VI) by Mg(OH)₂ in the presence of Al(III) in Section 2.3. Comparing the first adsorption, it was found that the efficiency of the second equilibrium adsorption increased significantly (from 40% to 74%), considering the LDH in the solid sample Mg(OH)₂.

3.3. The Mechanism of the Cr(VI) Adsorption

In this study, in order to explore the stability of various Cr(VI) adsorbed systems, Gauss View 6.0 was used to model the individual structures in this experiment, and the DFT method, a high-precision quantum mechanical approach, was used to geometrically optimize the structures of the adsorption systems at the B3LYPD3 level using Gaussian 09 software [36]. To calculate the energy, the 6-311G** basis set was used for H, Mg, and O, and the SDD basis set was used for Cr and Al. In addition, the Cr(VI) adsorption energy on the surface of Mg(OH)₂ can be computed from the following equation:

E_ads = E_{Mg(II)-Al(III)-oxyhydroxide} − E_Mg(II) − E_{Al(III)-oxyhydroxide}

(9)

where E_Mg(II) and E_{Al(III)-oxyhydroxide} represent the system’s energy following geometric optimization of Mg(II) and Al(III)-oxyhydroxide prior to adsorption, respectively. The theoretical analysis of the adsorption energy calculation was used to explain the experimental phenomena of first adsorption and then desorption of Cr(VI) by Mg(OH)₂ in the presence of Al(III). The adsorption model structure 1 is a simulation of Cr(VI) adsorption onto Mg(OH)₂ in the absence of Al(III), with E_ads = 27.9 kJ/mol (E_ads,1 > 0), showing that the process was a heat absorption reaction, which is more difficult to occur. In addition, the adsorption system in the presence of Al(III) was further analyzed by considering two binding forms of Al(III) complexed with HCrO₄⁻ (i.e., 2. Mg(OH)₂-HCrO₄⁻ + Al³⁺ and 3. Mg(OH)₂-Al³⁺-HCrO₄⁻) to simulate the possible adsorption structures (E_ads,2 = 69.1 and E_ads,3 = −8.8 kJ/mol). Where E_ads,3 < 0, the reaction was an exothermic process and can proceed spontaneously. Furthermore, in the case of the ternary complex structure 3. Mg(OH)₂-Al³⁺-HCrO₄⁻, Al(III) can quickly enter the lattice and form the structure 4. LDH. As E_ads,4 > 0 (E_ads,4 = 40.6 kJ/mol), the adsorption reaction was difficult to proceed spontaneously, which was considered to be the cause of the desorption of Cr(VI). As an important element in heavy metal adsorption, pH influences the molecular potential on the adsorbent surface [37]. In this system, the increase in pH (Figure 4a) demonstrated that OH⁻ bound to Al(III) and competed with Cr(VI) for the adsorption sites on Al(III). The complexes of Cr(VI) with Al(III) were readily transformed back to the free state, improving the possibility of their desorption. However, as this system proceeded, the surface of LDHNSs became positively charged, leading to rapid electrostatic adsorption of negatively charged Cr(VI) in the proximity of LDHNSs. Since the Cr(VI) was intercalated into the formed LDH, it was more immobile and difficult to desorb.

Based on the fact that the above processes occurred simultaneously, it was concluded that the solution containing Cr(VI) and Al(III) had an initial pH of 4 in this experiment. With the addition of Mg(OH)₂ adsorbent, the positive charge of Al(III) was preferred to adsorb on the surface of Mg(OH)₂ (Isoelectric Points 11.9) [17] through a chemical binding reaction, and two types of ternary surface complexes might be formed (Equations (5) and (6)). Therefore, the promotion of Cr(VI) adsorption efficiency by ternary complexes in the presence of Al(III) was accompanied by the replacement process of the aluminum lattice. With the increase in pH in the reaction, OH⁻ occupied surface sites [38], and the ternary surface complex that preferentially absorbed HCrO₄⁻ onto Mg(OH)₂ could be released from the surface. However, due to the electrostatic adsorption of the LDHNSs layers formed after lattice replacement, the desorbed HCrO₄⁻ was more immobile between the LDH layers. And this phenomenon that the competitive adsorption of OH⁻ led to partial Cr(VI) desorption was consistent with the experimental results in Figure 1a. Therefore, it was possible that the reason for the desorption of Cr(VI) in this experiment could be that Cr(VI) was not adsorbed by the formed LDH.

To explore the impact of the cation in the reaction, 1 mM KCl was added to the solution containing Cr(VI) and Mg(OH)₂. In this experiment, the results of Cr(VI) removal did not make any significant difference without the addition of Al(III) ions. K(I) did not promote the removal of Cr(VI) or the formation of LDH, which was different from the experiments in the presence of Al(III). Furthermore, different aluminum salts were added to study the influence of anions. The addition of nitric acid and chloride ions led to consistent results showing that the adsorption of Cr(VI) had been promoted (Figure 4b). However, the addition of sulfate was not shown to facilitate the adsorption process of Cr(VI), and the reaction was essentially the same as when no other ions were added. This was probably because sulfate and chromate are competing for adsorption on the solid surface. It was also demonstrated in the literature [39] that anionic sulfates with higher valence states are more likely to enter the interlayer when forming LDH. Therefore, the competitive adsorption of LDH in the intercalation layer was also considered the main reason for the poor removal of Cr(VI). In conclusion, anions with the same negative charge and lower valence had no influence on Cr(VI) adsorption, while high-valent sulfates had a greater effect than lower-valence anions due to competitive adsorption on the surface of Mg(OH)₂ and in the interlayer of the formed LDH.

In summary, the kinetics of Cr(VI) adsorption on Mg(OH)₂ in the presence of Al(III) were divided into three stages. Firstly, the ternary surface complex formation promoted the almost complete adsorption of Cr(VI) (As illustrated in Figure 1a, the maximum adsorption efficiency of Cr(VI) was 91%). Second, the weak surface adsorption led to Cr(VI) desorption because of the variation in pH values. Third, the interlayer adsorption of Cr(VI) formed stable LDH (see the schematic diagram of the LDH formation mechanism in Figure 5). Thus, the reaction mechanism of Cr(VI) adsorption is that at the beginning of the reaction, two ternary surface complexes were formed on the Mg(OH)₂ absorbed Cr(VI) (Equations (5) and (6)). As the reaction progressed, the cationic (Al(III)) isomorphic substitution and the formation of LDH led to the anionic adsorption of Cr(VI) in the gallery (Equation (7)). Therefore, it could be concluded from the above analysis that the Cr(VI) complexes formed on the Mg(OH)₂ surface are weak and susceptible to changes in environmental behavior (pH), leading to Cr(VI) desorption. However, after the formation of hydrotalcite, chromate became more immobile and difficult to release under neutral and alkaline conditions.

4. Conclusions

In this paper, the adsorption of Cr(VI) to Mg(OH)₂ minerals in the presence of Al(III) was investigated. The results show that the presence of Al(III) not only enhanced the adsorption and removal of Cr(VI) by Mg(OH)₂ nanosheets but also formed the Mg(II)-Al(III)-LDH, which facilitated the immobilization of Cr(VI). However, the change in pH during the reaction led to competitive adsorption and caused the desorption of Cr(VI). Therefore, based on the fact that the actual complex chromium-containing industrial wastewater often contains other pollutants, this study could provide certain guidance and references for the methods of co-adsorption, removal, and purification of complex pollutants using magnesium hydroxide or other minerals. Meanwhile, studying the environmental migration of Cr(VI) in Mg(OH)₂ minerals could also improve our knowledge of the environmental pollutants migration mechanisms in geochemically complex systems.