**Synthesis of MgFe Layered Double Hydroxide from Iron-Containing Acidic Residual Solution and Its Adsorption Performance**

#### **Alin Golban, Lavinia Lupa \*, Laura Cocheci \* and Rodica Pode**

Politehnica University of Timisoara, Faculty of Industrial Chemistry and Environmental Engineering,

6 Vasile Parvan Blvd., Timisoara 300223, Romania; rodica.pode@upt.ro

**\*** Correspondence: lavinia.lupa@upt.ro; laura.cocheci@upt.ro; Tel.: +402-5640-4159

Received: 19 August 2019; Accepted: 1 October 2019; Published: 3 October 2019

**Abstract:** The paper presents a new method of layered double hydroxide (LDH) synthesis starting from secondary sources, namely acidic residual solutions. The iron content of the acidic solution resulting from the pickling step of the hot-dip galvanizing process make it suitable to be used as an iron precursor in LDH synthesis. Here, Mg4Fe–LDH synthesized through the newly proposed method presented structural and morphological characteristics similar to the properties of layered double hydroxides synthesized from analytical-grade reagents. Moreover, the as-synthesized LDH and its calcined product presented efficient adsorption properties in the removal process of Mo(VI) from aqueous solutions. The adsorption studies are discussed from the equilibrium, kinetic, and thermodynamic points of view. The proposed novel technologies present both economic and environmental protection benefits.

**Keywords:** iron precursor; acidic residual solution; LDH synthesis; Mo(VI) adsorption

#### **1. Introduction**

Concerns over water resource pollution are continuously increasing, which has intensified research efforts regarding water decontamination. Because issues arise when the contaminants are in trace amounts but still at a concentration which exceeds the maximum admitted values, one of the most studied water treatment methods is adsorption. The adsorption process has gained researchers' interest due to its simple operating conditions and the versatile types of adsorbent materials which exist on the market [1–3]. Among the multitude of adsorbent materials, a considerable amount of attention has been paid to layered double hydroxide (LDH) compounds. The general formula of an LDH is [MII1 <sup>−</sup> xMIIIx (OH)2] <sup>x</sup>+. [An−x/<sup>n</sup> . mH2O]<sup>x</sup>−, where MII is a divalent cation, MIII is a trivalent cation, and An<sup>−</sup> is an anion. Their lamellar structure is based on brucite-like sheets, where some divalent cations are replaced with trivalent cations, resulting in some positively charged layers that contain between them various anions such as CO3 <sup>2</sup>−, Cl<sup>−</sup>, NO3 −, or even organic anions [4,5]. This structure increases their adsorptive properties, especially if the contaminant is in the form of oxyanions [6–8]. Due to the positive charge of the brucite-like lamellar layers, one of the main properties of synthetized LDHs is anionic exchange. If the LDH is thermal treated at temperatures up to 600 ◦C, when the obtained mixed metal oxides are immersed in aqueous solution, they are able to rehydrate and restore the lamellar structure of the LDH while retaining the anions present in solution in order to provide a neutral LDH molecule. This property is referred to as the "memory effect" and it is most often utilized to treat water containing undesirable anions or to introduce various anions into the LDH structure. From the multitude of Me<sup>2</sup>+/Me3<sup>+</sup> LDH types, much focus has been directed toward the Mg<sup>2</sup>+/Fe3<sup>+</sup> pair. A literature search revealed that Mg/Fe–LDH has been used for phosphate removal from aqueous solutions [9–11]; as adsorbent materials of various heavy metals, such as Cr, As, Pb, Zn, Cu, Se, Sb, and

so forth [12–16]; for treatment of aqueous solutions contaminated with different reactive dyes [17–19]; or as a catalyst [20–22].

Presently, there is a focus on obtaining cheaper adsorbent materials that still retain properties similar to those obtained from pure chemicals. Therefore, the tendency is to replace some raw materials with secondary sources of various metals. This helps reduce production costs as well as protect the environment by recycling and recovering various waste products. To this end, some researchers have studied the possibility of obtaining LDHs by using, as a precursor of various metal ions, different industrial wastes, such as fly ash, zinc ash, furnace slag, aluminum slag, and so forth. [23–27]. In most cases, the wastes were used as precursors for magnesium, zinc, or aluminum ions. However, even though the obtained LDHs presented very good properties similar to LDHs obtained from reagents and the process helped the environment by recycling wastes, there were no significant reductions in the production cost. In all of these cases, the precursor was a solid waste, which first needed to be brought into solution. This required the use of some acid solutions and added an extra step in the LDH production process.

Therefore, in the present study, we present a new Mg/Fe–LDH obtained from secondary sources using, for the first time, an acidic residual solution resulting from the pickling step in the hot-dip galvanizing process as an iron precursor. In a previous study, we demonstrated that an Mg/Fe–LDH obtained using a secondary source as an iron precursor (iron sludge resulting from the hot-dip galvanizing process) presented properties similar to those obtained from reagents [28]. The novelty of this method is that, by directly using an acidic waste solution as an iron precursor, two steps are avoided in the process of obtaining LDHs. First, the acidic solution from the hot-dip galvanizing process is not neutralized in order to obtain the sludge, and second, the hydrometallurgical leaching of sludge is avoided when obtaining the iron precursor. In order to determine the efficiency of using the acidic residual solution as an iron precursor in the process of obtaining Mg/Fe–LDH, the LDH was analyzed and used as an adsorbent material in the removal process of Mo(VI) as molybdate (MO4 <sup>2</sup>−) from aqueous solutions.

Due to the extensive use of Mo(VI) in many practices, molybdate anions (MO4 2–), the most common oxyanions of Mo(VI), can be found in various waste waters, such as mining waters, scrubber effluent of municipal solid waste incinerators, waste waters from the production of stainless-steel or cast iron alloys, and waste water from the production of various pigments and catalysts for high-temperature chemical processes [29–31]. If molybdenum is present in drinking water at a level higher than 0.07 mg/L, which is the maximum value admitted by the World Health Organization (WHO) [32], it could be toxic to animals and humans [33]. Therefore, finding an efficient method to remove oxyanions from aqueous solutions is still a challenge, taking into account the fact that the traditional method of waste water treatment, precipitation, can remove only cations and not oxyanions [30]. There are few papers which report the removal of Mo(VI) from aqueous solutions involving an ionic exchange mechanism using various LDHs synthesized from pure reagents [7,8,34]. The purpose of this study was to compare the adsorption performance of the new, synthesized LDH from a secondary iron source with those reported in the literature.

#### **2. Experimental**

#### *2.1. Mg*/*Fe–LDH Synthesis and Characterization*

The acidic residual solution was received from a local hot-dip galvanizing plant and was subjected to chemical analysis in order to determine the metal ion concentrations. The concentration of Fe2<sup>+</sup> was determined through titration with KMnO4. The concentrations of total iron and other ions present in solution were determined by atomic absorption spectrometry using a Varian SpectrAA 280 FS spectrophotometer (Agilent, Santa Clara, CA, USA).

Mg/Fe–LDH was synthesized using the coprecipitation method at low oversaturation [4]: 200 mL of 1 M solution containing iron residual solution and magnesium nitrate, at a Mg:Fe molar ratio of 4:1, was added dropwise under continuous stirring to 100 mL of 1 M Na2CO3 solution. The pH of the suspension was maintained in the range of 10–11 using a 2 M NaOH solution. The suspension was aged at 70 ◦C for 20 h, then filtered and washed with distilled water several times until reaching a pH of 7. The obtained slurry was dried overnight, then milled and sieved in order to obtain particles with dimensions less than 90 μm, which were used in this study. A portion of the obtained Mg4Fe–LDH was calcined at 450 ◦C at a heating rate of 10 ◦C/min. It was maintained at this temperature for 3 h using a Nabertherm oven. The obtained material was named Mg4Fe-450.

The as-synthesized sample and the calcined one were characterized by powder X-ray diffraction (XRD), scanning electron microscopy (SEM), X-ray dispersion analysis (EDX), and specific surface area and pore volume. The RX diffractograms were recorded using a Rigaku Ultima IV X-ray diffractometer (Rigaku Analytical Devices Inc., Wilmington, MA, USA). SEM images were recorded using a Quanta FEG 250 microscope (FEI Company, Hillsboro, OR, USA) equipped with a ZAF-type EDX quantifier (FEI Company). A Micromeritics ASAP 2020 instrument (Micromeritics, Norcross, GA, USA) was used to determine the specific surface area and pore volume of the studied samples.

In order to determine the adsorption performance of the obtained materials, both samples—the as-synthesized LDH (Mg4Fe) and the calcined one (Mg4Fe-450)—were used in experiments to remove Mo(VI) from aqueous solutions.

The resulting adsorbent after Mo(VI) adsorption was analyzed using an FEI Tecnai F20 G2 TWIN TEM (FEI Company) at an accelerating voltage of 200 kV in bright field mode in order to determine the morphology.

#### *2.2. Adsorption Studies*

Both materials, the as-synthesized and the calcined samples, were used in the removal process of Mo(VI) anions from aqueous solutions. The adsorption process was conducted in batch mode using a Julabo SW23 shaker for sample shaking at a constant rotation speed (200 rpm). The adsorption performance of the obtained materials developed in the removal process of Mo(VI) anions from aqueous solutions was studied taking into account the influence of the Mo(VI) initial concentrations (15–200 mg/L), stirring time (5–240 min), and temperature (20, 35, and 50 ◦C). For all the adsorption studies, the solid:liquid ratio between the adsorbent and Mo(VI)-containing aqueous solutions was 0.025 g in 25 mL (i.e., solid:liquid ratio (S:L) = 1 g/L). The initial pH of the solution was adjusted to pH = 6.0 ± 0.5. According to the literature review and our previous studies [2,35], Mo(VI) adsorption onto LDHs is maximal at this pH due to the fact that, at this value, Mo(VI) is found in the solution under MoO4 2−.

After each adsorption experiment, the samples were filtered in order to separate the phases and to analyze the residual concentration of Mo(VI). The Mo(VI) concentration before and after adsorption was determined through atomic absorption spectrometry using a Varian SpectrAA 280 FS spectrometer.

The mass balance presented in Equation (1) was used for the calculation of Mo(VI) uptake on 1 g of adsorbent:

$$\mathbf{q}\_{\mathbf{c}} = \frac{(\mathbf{C}\_0 - \mathbf{C}\_{\mathbf{c}}) \cdot \mathbf{V}}{\mathbf{m}} \tag{1}$$

where qe is the adsorbed quantity of Mo(VI), expressed in milligrams per 1 g of studied adsorbent; Co and Ce represent the Mo(VI) concentration of the aqueous solutions before adsorption and after the established equilibrium (mg/L); V is the volume of the Mo(VI)-containing aqueous solution (L); and m is the mass of the Mg4Fe–LDH (g) used in the experiments.

#### **3. Results and Discussions**

#### *3.1. Material Characterizations*

It is estimated that, in the European Union, almost 300,000 m3 of acidic residual solution is produced annually in the hot-dip galvanizing industry. [36] Due to its corrosive nature and high concentration of metal ions, this solution is considered a toxic waste and requires treatment before dumping. The treatment processes for spent pickling solutions can be classified as neutralization treatments, treatments for acid recovery, and treatments for metal recovery. Neutralization treatments consist of adding NaOH or Ca(OH)2 when the metal ions are precipitated, and the resulting sludge is dried and considered a solid waste. Treatment processes for acid recovery include evaporation, membrane processes (membrane dialysis, membrane electrolysis, and membrane distillation), and processes with fluidized beds (Ruthner or spray-roasting processes). Metal ion recovery can employ ion exchange, crystallization, or solvent extraction. The physicochemical characterization of the acidic residual solution resulting from the pickling step of the hot-dip galvanizing process is presented in Table 1. It can be observed that the residual solution contains a high concentration of iron (65 g/L), which makes it suitable to be used as an iron precursor in LDH synthesis. The concentration of iron ions in the residual solutions can be considered constant because these acidic solutions resulting from the pickling step in the hot-dip galvanizing industry are continuously analyzed, and when they achieve the maximum value of iron ions and the minimum value of acidity, they are removed from the process and can be used as a secondary source of iron ions for LDH manufacturing applications. Besides iron ions, the residual solution also contains other metal ions, but in a smaller concentration; their presence does not inhibit the synthesis of LDHs.


**Table 1.** Physicochemical characterization of the acidic residual solution.

The XRD patterns of the obtained adsorbent are presented in Figure 1. It can be observed that the Mg4Fe obtained from a secondary source of iron showed a unique main phase which corresponded to pyroaurite. The cell parameters a and c were calculated using Bragg's law, assuming a rhombohedral symmetry of crystallization. The cell parameter a represents the cation–cation distance within the brucite-like layer (a = 2 . d(110) = 3.11 Å) and c represents the adjacent distance of the hydroxide layer (c = 3 . d(003) = 23.9 Å), where d = λ/2 sin θ and λ = 1.54056 Å. The obtained results are in good agreement with other results reported by several authors [19,28,37]. The impurities present in the acidic residual solution were under 5% and they did not lead to the formation of secondary phases in the Mg4Fe–LDH structure. In the case of the sample calcined at 450 ◦C, the LDH structure was modified, and the RX diffractogram was specific to a poor crystalline phase corresponding to MgO (periclase), with Fe3<sup>+</sup> probably dispersed in the structure. Due to smaller Fe3<sup>+</sup> than Mg2<sup>+</sup> concentration utilized in the starting solution and the low temperature of the thermal treatment (450 ◦C) of Mg4Fe, it was difficult to achieve the formation of iron oxide species such as magnesioferrite (MgFe2O4) or maghemite (Fe2O3). At this calcination temperature, the iron oxides were amorphous and the crystals were in the course of forming; therefore, they could not be identified by the RX analysis. These findings are also in agreement with other results reported in the literature [37,38]. For this reason, it is expected that this sample has the highest adsorption capacity in the removal process of Mo(VI) anions from aqueous solutions [19,28].

**Figure 1.** XRD pattern of the synthesized materials.

The morphology of the synthesized samples can be observed from the SEM images presented in Figure 2. The Mg4Fe–LDH presented an aerated structure of fluffy particles. Through calcination, the sample became more amorphous, and the surface had aspects of cotton flowers.

**Figure 2.** SEM images and X-ray dispersion analysis (EDX) spectra of the synthesized samples: (**a**) Mg4Fe and (**b**) Mg4Fe–450.

The EDX spectra presented in Figure 2 show the peaks of the characteristic elements of the studied materials. No characteristic peaks appeared for the impurities present in the precursor solution. This indicates that there was a negligible quantity of impurities.

The molar ratio between Mg and Fe ions from the studied samples together with their BET specific surface area and pore volume are presented in Table 2.


**Table 2.** Physical and chemical properties of the studied samples.

The impurities present in the spent pickling solution were analyzed in the washing solutions of the LDH in the as-synthesized solid samples and the calcined one and in the solutions after Mo(VI) adsorptions. It can be observed from the chemical analysis that impurities could be found in the solid synthesized samples but in a concentration under 2%, this being the reason why these do not appear in the EDX and RX analyses, as they are under the detection limit of 5%. Impurities could not be detected in the washing waters of the synthesized Mg4Fe–LDH and in the residual solutions after Mo(VI) adsorption. This demonstrates that the impurities were not released during the adsorption process from the solid support.

#### *3.2. Equilibrium Studies*

The equilibrium isotherms representing the dependence of the adsorption capacities developed by the Mg4Fe and Mg4Fe-450 as a function of the Mo(VI) concentrations at equilibrium are presented in Figure 3. Increasing the initial concentration of Mo(VI) increased the active sites available for adsorption and, therefore, increased the adsorption capacities of both the studied materials. Mg4Fe developed an experimental maximum adsorption capacity in the removal process of Mo(VI) from aqueous solutions of 39.9 mg/g. Further, due to the memory effect, the calcined samples exhibited a higher experimental maximum adsorption capacity of almost 50% (qe = 52.8 mg/g). For aqueous solutions with Mo(VI) initial concentrations of ≤30 mg/L, the removal degree of Mo using a S:L ratio of 1 g/L was higher than 90%. For aqueous solutions with higher Mo(VI) initial concentrations, in order to obtain higher removal degrees, it was necessary to increase the S:L ratio. Other studies reported in the literature also present higher adsorption capacities for calcined samples compared with synthesized samples [18,28].

**Figure 3.** Equilibrium isotherms of Mo(VI) adsorption onto Mg4Fe and Mg4Fe-450. S:L = 1 g/L, t = 60 min, pH = 6, T = 20 ◦C.

Every adsorption study aims to determine the maximum adsorption capacity of the studied adsorbents and the equilibrium coefficient in order to achieve the design. Therefore, several isotherms, such as Langmuir, Freundlich, Temkin, and Dubinin–Radushkevich (DR) isotherms in their linear form, have been employed for this purpose [2,18,39]. The Langmuir isotherm supposes that the adsorption process takes place in a single layer on a uniform surface containing equivalent sites of the studied adsorbents. The Freundlich isotherm is used to express the affinity of the studied adsorbent to the retained pollutant. If the adsorption data fit the Temkin isotherm well, this means that the adsorbent surface is heterogeneous. In order to design the equilibrium adsorption regarding Mo(VI) removal using Mg4Fe and Mg4Fe-450, the linear graphs of the mentioned isotherms were plotted (Figure 4) and the obtained equilibrium isotherm parameters together with the regression coefficients are presented in Table 3.

**Figure 4.** Equilibrium isotherms of Mo(VI) adsorption onto the studied adsorbent: (**a**) Langmuir, (**b**) Freundlich, (**c**) Temkin, and (**d**) Dubinin–Radushkevich.


**Table 3.** Equilibrium sorption isotherm parameters for Mo(VI) adsorption onto Mg4Fe and Mg4Fe-450.

Both studied materials presented affinity for Mo(VI) ion removal (1/n parameters were below unity).

By comparing the equilibrium isotherm parameters presented in Table 3, it can be concluded that Mo(VI) removal from the aqueous solutions occurred as a monolayer at the uniform surfaces of the Mg4Fe and Mg4Fe-450 materials because the Langmuir isotherm obtained the highest regression coefficients (closed to unity). Further, there was no significant difference between the maximum adsorption capacity calculated from its plot and those determined experimentally.

The lowest regression coefficients were obtained for the DR plots. At the same time, different values were obtained for the monomolecular adsorption capacities qs for the studied materials and the values experimentally obtained in the adsorption process of Mo(VI) removal from aqueous solutions. The results suggest that the DR plot cannot be used to model the adsorption data of MoO4 <sup>2</sup><sup>−</sup> onto Mg4Fe and Mg4Fe-450 [18,39].

#### *3.3. Kinetic Studies*

Kinetic studies were used to determine the optimum time necessary to establish the equilibrium between the Mg4Fe–LDH and the Mo(VI) anions. The experiments regarding the various stirring times were conducted at three different temperatures (Figure 5). Mo(VI) removal occurred quite quickly in the first minutes of contact between the adsorbent and adsorbate, especially when the calcined sample was used. After 60 min of stirring, the adsorption capacity increased slowly, so it can be considered that equilibrium was achieved in 60 min at all the studied temperatures for both adsorbent materials. The temperature increase led to a slight increase of the adsorption capacity of Mg4Fe and Mg4Fe-450 in the removal process of Mo(VI).

The well-known models of pseudo-first-order, pseudo-second-order, and intraparticle diffusion were used to simulate the kinetics of Mo(VI) adsorption onto the studied materials [18,19,40]. The representations of their linear plots are presented in Figures 6–8, and the kinetic parameters together with the obtained correlation coefficients are summarized in Table 4. From the linear representation of the experimental data according to the pseudo-first-order and pseudo-second-order kinetic models (Figures 6 and 7) and the interpretation of their parameters (Table 4), it can be concluded that Mo(VI) adsorption onto Mg4Fe and Mg4Fe-450, separately, is best described by the pseudo-second-order kinetic model. In this case, for all three temperatures, correlation coefficients were obtained that were close to 1, and the calculated adsorption capacities were similar to those determined experimentally.

**Figure 5.** Effect of contact time on the adsorption capacity of (**a**) Mg4Fe and (**b**) Mg4Fe-450 at three different temperatures in the removal process of MoO4 <sup>2</sup><sup>−</sup> from aqueous solutions. S:L = 1 g/L, C0 = 50 mg/L Mo(VI), pH = 6.

**Figure 6.** Pseudo-first-order kinetic model of MoO4 <sup>2</sup><sup>−</sup> adsorption onto (**a**) Mg4Fe and (**b**) Mg4Fe-450.

**Figure 7.** Pseudo-second-order kinetic model of MoO4 <sup>2</sup><sup>−</sup> adsorption onto (**a**) Mg4Fe and (**b**) Mg4Fe-450.

**Figure 8.** Intraparticle diffusion model of MoO4 <sup>2</sup><sup>−</sup> adsorption onto (**a**) Mg4Fe and (**b**) Mg4Fe-450.


**Table 4.** Kinetic model parameters for Mo(VI) adsorption onto Mg4Fe and Mg4Fe-450.

The pseudo-second-order kinetic model can be used to simulate the experimental data regarding MoO4 <sup>2</sup><sup>−</sup> adsorption onto the studied materials; this means that the process is controlled by a chemical sorption [19,40,41].

Through the simulation of the experimental data according to the intraparticle diffusion model (Figure 8), it can be observed that the straight line does not pass through origin, indicating that the rate-limiting step for Mo(VI) adsorption is not the intraparticle diffusion. Also, the straight line presents a fragmentation after a while, suggesting that the adsorption process is more complex. The fast removal of Mo(VI) in the first minutes of contact between the adsorbent and adsorbate was controlled by the film diffusion when the studied adsorbent surfaces were covered with MoO4 <sup>2</sup><sup>−</sup> anions. After the surface coverage, the transportation of Mo(VI) inside the adsorbent particles occurred, as suggested by the second straight line obtained through the representation of q versus t1/<sup>2</sup> (Figure 8) [19].

#### *3.4. Thermodynamic Studies*

From the linear plot representation of the Arrhenius equation (ln(K2) versus 1/T, Figure 9), the activation energy of Mo(VI) adsorption onto the studied materials was determined. The activation energy value confers information regarding the type of adsorption process (physical, when E < 4.2 kJ/mol, or chemical, when E > 4.2 kJ/mol) [33]. The activation energy determined for MoO4 2− sorption onto Mg4Fe and M4Fe-450 was determined to be 6.37 and 7.89 kJ/mol, respectively, indicating chemisorption (Table 5).

**Figure 9.** Arrhenius plot of MoO4 <sup>2</sup><sup>−</sup> adsorption onto the studied materials.


**Table 5.** Thermodynamic parameters for Mo(VI) adsorption onto Mg4Fe and Mg4Fe-450.

The thermodynamic equilibrium constant Kd, defined as the ratio between the Mo(VI) concentration in the solid phase and the solution at equilibrium, was used to calculate the Gibbs free energy (ΔGo) at the studied temperatures and for the van 't Hoff representation in order to determine the thermodynamic parameters enthalpy (ΔHo) and entropy (ΔSo) for MoO4 <sup>2</sup><sup>−</sup> adsorption onto the studied materials [19,40,41]. The plot of the van 't Hoff representation is shown in Figure 10 and the thermodynamic parameters are listed in Table 5.

**Figure 10.** Van 't Hoff plot of MoO4 <sup>2</sup><sup>−</sup> adsorption onto the studied materials.

The MoO4 <sup>2</sup><sup>−</sup> adsorption onto the Mg4Fe and Mg4Fe-450 materials was an endothermic process, due to the obtained positive value for ΔHo, and spontaneous because the Gibbs free energy presented negative values, which decreased as the temperature increased. Also, the positive values for ΔHo and ΔS<sup>o</sup> suggest that, during the adsorption process, the solid–liquid interface increases the randomness and the adsorption takes place due to chemical interactions [41].

#### *3.5. Adsorption Performance*

The adsorption performances of the synthesized materials developed in the removal process of MoO4 <sup>2</sup><sup>−</sup> adsorption from aqueous solutions were compared with the adsorption capacities developed by similar materials that were reported in the literature. The results are presented in Table 6, and it can be observed that Mg/Fe–LDH obtained from secondary sources can be efficiently used as an adsorbent material for treatment of water with dissolved Mo(VI).


**Table 6.** The comparison between the adsorption capacities of similar adsorbents developed for the treatment processes of aqueous solutions containing MoO4 <sup>2</sup><sup>−</sup> anions.

After adsorption, the solid materials were recovered by filtration and were subjected to XRD analysis. The XRD patterns of the Mg4Fe and Mg4Fe-450 are presented in Figure 11.

The as-synthesized material, Mg4Fe, did not suffer any change in its crystalline structure after adsorption, with the unit cell parameters being a = 3.11 Å and c = 24.0 Å (compared to the unit cell parameters before adsorption: a = 3.11 Å and c = 23.9 Å). The 003 reflection was almost the same for Mg4Fe before adsorption (d(003) = 7.967 Å) and after adsorption (d(003) = 8.000 Å). This demonstrates that the adsorption of Mo(VI), as molybdate anions, was performed on the LDH surface and not through anion exchange, considering that carbonate is the most difficult to replace among the anions present in the LDH interlayer gallery. On the other hand, as shown in Figure 11, the calcined material, Mg4Fe-450, regained its layered double hydroxide structure, and all the characteristic peaks of pyroaurite were present in the diffractogram. The unit cell parameters of Mg4Fe-450 after adsorption (a = 3.12 Å and c = 24.7 Å) suggest that the rehydration process of the calcined material (the memory effect) was developed

through adsorption of Mo(VI) as molybdate anions from the solution and incorporated in the interlayer space of LDH. Furthermore, d(003) for Mg4Fe-450 after adsorption increased to 8.233 Å. This increase in basal spacing indicates that molybdate anions were intercalated into the interlayer spaces of the reformed LDH. During Mo(VI) adsorption onto the calcined sample, instead of carbonate ions from the atmosphere, Mo(VI) anions were intercalated from the solutions due to the fact that the anion volume of Mo(VI) is higher than that of carbonate (VMoO4<sup>2</sup><sup>−</sup> = 0.088 nm3; VCO32<sup>−</sup> = 0.061 nm3) [43]. The retention of Mo(VI) through the memory effect between the interlayer gallery and also through adsorption onto the layer of the Mg4Fe-450 surface explain the higher adsorption capacity developed by this material compared with its precursor. These results agree with other results reported in the literature [34,44].

**Figure 11.** XRD patterns of the materials after Mo(VI) adsorption.

The morphology of the samples after Mo(VI) adsorption can be observed from the TEM images presented in Figure 12. The specific hexagonal morphology with ultrathin layers for the LDH samples can be observed in the presented TEM images. Due to their small crystallite size and the fact that it is difficult to isolate a unique crystal, the formation of some aggregates can be observed. The calcined sample, through the memory effect after Mo(VI) adsorption, returned to the hexagonal shape specific to LDH due to reformation of the layer structures.

**Figure 12.** TEM images of the materials after Mo(VI) adsorption: (**a**) Mg4Fe and (**b**) Mg4Fe-450.

#### **4. Conclusions**

This paper reports the successful synthesis of a new Mg4Fe–LDH using, as iron precursor, a secondary source, namely, the acidic residual solution resulting from the pickling step of the hot-dip galvanizing process. The obtained LDH presented similar properties to those obtained from pure reagents. The impurities present in the residual solutions, besides the iron ions, did not interfere with the structure of the obtained LDH. The obtained Mg4Fe and the calcined product presented efficient adsorption properties in the removal process of MoO4 <sup>2</sup><sup>−</sup> from aqueous solutions. The proposed method of obtaining LDH presents multiple benefits: (1) it decreases the cost of obtaining LDH; (2) it decreases the cost of acidic residual solution neutralization; (3) it reduces waste discharge into the environment, and (4) it minimizes the use of raw materials.

**Author Contributions:** Conceptualization, L.L. and R.P.; methodology, L.C.; formal analysis, A.G. and L.L.; investigation, A.G., L.L. and L.C.; writing—original draft preparation, L.L. and R.P.; writing—review and editing, L.L. and L.C.; visualization, L.C.; supervision, R.P.

**Funding:** This research received no external funding.

**Acknowledgments:** The studies were done during the PhD program from the Doctoral School of the University Politehnica Timisoara.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


© 2019 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (http://creativecommons.org/licenses/by/4.0/).

## *Article* **Antimonate Removal from Polluted Mining Water by Calcined Layered Double Hydroxides**

#### **Elisabetta Dore \*, Franco Frau and Rosa Cidu**

Department of Chemical and Geological Sciences, University of Cagliari, 09042 Monserrato, Cagliari, Italy **\*** Correspondence: elisabettadore@yahoo.it

Received: 20 June 2019; Accepted: 1 August 2019; Published: 6 August 2019

**Abstract:** Calcined layered double hydroxides (LDHs) can be used to remove Sb(V), in the Sb(OH)6 − form, from aqueous solutions. Sorption batch experiments showed that the mixed MgAlFe oxides, obtained from calcined hydrotalcite-like compound (3HT-cal), removed Sb(OH)6 − through the formation of a non-LDH brandholzite-like compound, whereas the mixed ZnAl oxides, resulting from calcined zaccagnaite-like compound (2ZC-cal), trapped Sb(OH)6 − in the interlayer during the formation of a Sb(V)-bearing LDH (the zincalstibite-like compound). The competition effect of coexistent anions on Sb(OH)6 <sup>−</sup> removal was HAsO4 <sup>2</sup><sup>−</sup> >> HCO3 <sup>−</sup> ≥ SO4 <sup>2</sup><sup>−</sup> for 2ZC-cal and HAsO4 2− >> HCO3 <sup>−</sup> >> SO4 <sup>2</sup><sup>−</sup> for 3HT-cal. Considering the importance of assessing the practical use of calcined LDHs, batch experiments were also carried out with a slag drainage affected by serious Sb(V) pollution (Sb = 9900 μg/L) sampled at the abandoned Su Suergiu mine (Sardinia, Italy). Results showed that, due to the complex chemical composition of the slag drainage, dissolved Sb(OH)6 − was removed by intercalation in the interlayer of carbonate LDHs rather than through the formation of brandholzite-like or zincalstibite-like compounds. Both 2ZC-cal and 3HT-cal efficiently removed very high percentages (up to 90–99%) of Sb(V) from the Su Suergiu mine drainage, and thus can have a potential application for real polluted waters.

**Keywords:** layered double hydroxides; antimonate uptake; mine water; brandholzite; zincalstibite

#### **1. Introduction**

Antimony (Sb) is an element widely present in the environment as a result of both natural processes and anthropogenic sources [1]. Due to its potential risk for human health, the World Health Organization has set the guideline value for drinking water at 20 μg/L of Sb [2], while the European Community has established 5 μg/L [3]. The Sb concentration in uncontaminated freshwater is usually lower than drinking water limits [4,5], however considerable higher concentrations (up to mg/L) can be related to both natural sources and anthropogenic activities [1,4,6–8]. In natural environments, Sb is generally present in the trivalent Sb(III) and pentavalent Sb(V) oxidation states, with the Sb(III) species being ten times more toxic than the Sb(V) ones [9]. In aqueous solution, Sb(III) and Sb(V) prevail, respectively, under reducing and oxidizing conditions as antimonous acid H3SbO3 and antimonic acid H3SbO4 and their dissociation products, with the Sb(OH)6 − anion being the most common and stable aqueous species in a wide range of natural pH values [1,9].

Among the techniques suitable for the abatement of Sb concentration in the solution, such as coagulation-flocculation [10], electrochemical methods [11,12] and membrane separation [13,14], the adsorption is considered a low cost and effective method [9]. Several studies reported that metal hydroxides and oxohydroxides (e.g., MnOOH, Al(OH)3, FeOOH) are good Sb removers, however they result more efficient for Sb(III) than Sb(V) under slightly acid to acid conditions [15–18]. Also, nano-TiO2 electroactive carbon nanotube (CNT) filter and ZrO2-carbon nanofibers (ZNC) were tested, respectively, for Sb(III) and simultaneous Sb(III) and Sb(V) removal; in particular, it was reported that

the adsorption capacity of Sb(III) is much higher than Sb(V) in ZNC [19,20]. Recent works showed that the acidic conditions are also favorable for the Sb(V) removal from solution by La and Ce-doped magnetic biochars [21,22]. With respect to the other sorbents, layered double hydroxides (LDHs) show the advantage of being able to remove hazardous anions from solution at circumneutral pH, and can therefore be potential removers for Sb(OH)6 − at the circumneutral pH and oxidizing conditions usually found in the environment [23].

The LDHs are minerals with general formula [M2<sup>+</sup>1-xM3<sup>+</sup>x(OH)2](An<sup>−</sup>)x/n·mH2O, where M2<sup>+</sup> and M3<sup>+</sup> are respectively bivalent and trivalent metals (Mg2<sup>+</sup>, Zn2<sup>+</sup>, Ca2<sup>+</sup>, Al3<sup>+</sup>, Fe3<sup>+</sup>, etc.), An- are anions (Cl<sup>−</sup>, SO4 <sup>2</sup><sup>−</sup>, CO3 <sup>2</sup>−, etc.) and x is the M<sup>3</sup>+/(M2<sup>+</sup> <sup>+</sup> M3<sup>+</sup>) molar ratio (0.20 ≤×≤ 0.33). The LDHs structure consists of octahedral brucite-like layers positively charged due to the partial substitution of M2<sup>+</sup> by M3<sup>+</sup>, stacked along the *c* axis and intercalated with interlayer anions, which neutralize the positive charge, and variable quantity of water molecules [24,25]. The LDHs can successfully remove anionic contaminants from solution through the anion exchange with the interlayer anions, or by trapping anions in the interlayer region during the reconstruction of the layered structure by the rehydration of mixed metal oxides obtained from LDHs calcination (the so called "memory effect") [23,26–29].

Previous authors reported that LDHs with different compositions are potential Sb(OH)6 − removers: nitrate and chloride bearing Mg-Al LDHs, both untreated and calcined, efficiently remove Sb(OH)6 − from solution through the formation of a brandholzite-like compound [30,31], and the removal capacity can be improved by doping LDHs with Fe2<sup>+</sup> [32]; sulfate bearing Zn-Al and Zn-Fe(III) LDHs uptake Sb(OH)6 − from solution by anion exchange [33,34] and also Fe-Mn LDHs obtained by electro-coagulation process result good removers [35]. Although interest in the use of LDHs for Sb removal from aqueous solutions has increased in recent years, to the best of our knowledge their practical use with real polluted water has not been investigated yet. Therefore, the aim of this study was to assess the Sb(OH)6 − removal capacity of calcined LDHs from real water affected by serious Sb pollution.

In our previous work we showed that calcined synthetic LDHs with composition like hydrotalcite (with formula Mg6(Al0.5Fe0.5)2(CO3)(OH)16·4H2O) and zaccagnaite (with formula Zn4Al2(CO3)(OH)12·3H2O) remove Sb(OH)6 − from solution, respectively, through the formation of a brandholzite-like phase (a non-LDH mineral with general formula Mg[Sb(OH)6]2·6H2O) and a zincalstibite-like compound (an LDH mineral with general formula Zn2Al(OH)6[Sb(OH)6]) [36]. In this work we used calcined hydrotalcite-like and zaccagnaite-like compounds to carry out batch experiments with coexistent anions in solution to evaluate their competition effect on Sb(OH)6 − removal. We successively assessed the practical use of these sorbents with real water by sorption batch experiments performed with the drainage water flowing out from the foundry slag impoundments at the abandoned mine of Su Suergiu (Sardinia, Italy), which is affected by serious Sb pollution [6,37].

#### **2. Materials and Methods**

#### *2.1. LDHs Synthesis and Calcination*

Synthetic hydrotalcite Mg6(Al0.5Fe0.5)2(CO3)(OH)16·4H2O and zaccagnaite Zn4Al2(CO3)(OH)12·3H2O were prepared with a coprecipitation method at constant pH [36]. Depending on composition, a solution (0.2 M) with the desired metals was prepared by dissolving in ultrapure water (Millipore, Milli-Q©, 18.2 MΩ cm) appropriate amounts of Mg(NO3)2·6H2O, Al(NO3)3·9H2O, Fe(NO3)3·9H2O and Zn(NO3)2·6H2O. All reagents were of analytical grade (ACS-for analysis, CARLO ERBA Reagents S.r.l., Cornaredo (MI), Italy) and were used without further purification. The so-obtained metal solution was dropped into a reactor containing a Na2CO3 solution (0.05 M), under stirring (500 rpm), and the precipitation was induced at constant pH (ranging between 9.5 and 10.5) by adding dropwise a NaOH (0.5 M) solution. After 24 h of aging at 65 ◦C, the solids were recovered through filtration (30 μm pore size cellulose filter, Whatman Plc, Little Chalfont, Buckinghamshire, UK), washed with deionized water and dried at room temperature. Calcination was performed at 450 ◦C for 4 h. The M<sup>2</sup>+/M3<sup>+</sup> molar ratios of synthetic

hydrotalcite and zaccagnaite, and their calcined products, were close to those of the starting solutions (Table 1).

From now on, synthetic hydrotalcite is termed 3HT and zaccagnaite is 2ZC, the numbers before the labels indicate the M2+/M3<sup>+</sup> molar ratio; moreover, the suffix -CO3 will be used for the untreated carbonate LDHs and the suffix -cal for the calcined LDHs.

**Table 1.** Chemical composition of synthetic untreated carbonate (-CO3) and calcined (-cal) hydrotalcite-like (3HT) and zaccagnaite-like (2ZC) compounds.


#### *2.2. Sorption Experiments*

#### 2.2.1. Effect of Coexistent Anions

In batch experiments, the coexistent anions were selected taking into account the chemical composition of Su Suergiu mine drainage [6,37]. The experimental solutions were prepared dissolving appropriate amounts of KSb(OH)6 and Na2SO4, NaHCO3 or Na2HAsO4·7H2O (ACS-for analysis, CARLO ERBA Reagents S.r.l., Cornaredo (MI), Italy) in ultrapure water. To perform the experiments, 0.1 g of 2ZC-cal or 3HT-cal was suspended, for 48 h under stirring, in 400 mL of solution containing equal concentrations (about 1 mmol/L) of dissolved Sb(OH)6 − and one competitor at a time. Experiments with only dissolved Sb(OH)6 − without competitors were also carried out to compare the results. Before the addition of the sorbents and during the experiments the pH of solutions was monitored and a portion of solution was withdrawn and acidified with HNO3 1% *v*/*v* for chemical analysis of Sb, As, S, Mg, Zn, Al and Fe by inductively coupled plasma optical emission spectroscopy (ICP-OES, ARL Fisons 3520, Waltham, MA, USA). At the end of the reaction time the solids were recovered through filtration (0.45 μm pore size polycarbonate filters, Whatman Plc, Little Chalfont, Buckinghamshire, UK), washed with distilled water and dried at room temperature for mineralogical characterization.

#### 2.2.2. Sorption Experiments with Su Suergiu Mine Drainage

The real water for sorption experiments is a slag drainage sampled at the abandoned Su Suergiu mine (Sardinia, Italy), at the sampling point named SU1 (Supplementary Materials Figure S1) as reported by Cidu et al. [37]. The physical and chemical parameters were determined at the sampling site using the sampling protocol described in Cidu et al. [37]. After sampling, the slag drainage (from now on SU1) was stored in HDPE bottles at 4 ◦C, and batch experiments were carried out, at room temperature (25 ◦C), less than 24 h after the sampling.

Different amounts of 3HT-cal or 2ZC-cal, equal to 0.1, 0.25, 0.5 and 1 g, were suspended in 400 ml of SU1 under stirring for 24 h. During the experiments the pH of solutions was monitored. At the end of reaction time, the solids were separated from solution through filtration, washed with distilled water and dried at room conditions for mineralogical characterization. The solutions recovered after the experiments were stored in two different aliquots: one aliquot was unacidified for analysis of major ions by ion chromatography (IC, Dionex ICS3000, ThermoFisher SCIENTIFIC, Waltham, MA, USA); a second aliquot was acidified for trace elements analysis (Sb, As, Fe, Zn and Al) by inductively coupled plasma mass spectrometry (ICP-MS, quadrupole, PerkinElmer SCIEX ELAN DRC-e, Waltham, MA, USA) with Rh as internal standard; concentrations of Sb > 1000 μg/L were also determined by ICP-OES.

#### *2.3. Mineralogical Characterization*

Mineralogical characterization of synthetic LDHs and their calcined products before and after the experiments was performed by collecting XRD patterns in the 5–80◦ 2θ angular range on an automated Panalytical X'pert Pro diffractometer (PANalytical, Almelo, Netherlands), with Ni-filter Cu-Kα<sup>1</sup> radiation (λ = 1.54060 Å), operating at 40 kV and 40 mA, using the X'Celerator detector.

#### **3. Results**

#### *3.1. E*ff*ect of Coexistent Anions on Sb(V) Removal*

#### 3.1.1. Sorptive Competition

The results of competition experiments showed that 2ZC-cal was slightly more effective than 3HT-cal (Figure 1). At the end of the experiments without competitors, 2ZC-cal removed 85% of Sb(OH)6 <sup>−</sup> whereas 3HT-cal reached 72%. The coexistence of SO4 <sup>2</sup><sup>−</sup> and HCO3 − slightly affected the Sb(OH)6 <sup>−</sup> uptake by 2ZC-cal, whereas in the experiments with 3HT-cal, the percentage of Sb(OH)6 − removed did not vary significantly in the presence of SO4 <sup>2</sup><sup>−</sup> but decreased up to 50% with coexistent HCO3 − (Figure 1, Tables 2 and 3).

**Figure 1.** Percentage of Sb(OH)6 − removed from solution by 2ZC-cal and 3HT-cal at the end of the experiments without competitors (Sb) and with HCO3 <sup>−</sup> (Sb/HCO3), SO4 <sup>2</sup><sup>−</sup> (Sb/SO4) or HAsO4 2− (Sb/HAsO4) as coexistent anions.

The HAsO4 <sup>2</sup><sup>−</sup> anion was the strongest competitor in the experiments with both 2ZC-cal and 3HT-cal, with percentages of Sb(OH)6 <sup>−</sup> removed lower than 10%. Moreover, the HAsO4 <sup>2</sup><sup>−</sup> concentration at the end of the experiment with 3HT-cal markedly decreased by about 60% (Table 3), showing a strong affinity of HAsO4 <sup>2</sup><sup>−</sup> for the interlayer region of 3HT.


**Table 2.** Solution pH values and dissolved ions determined before and at the end of the sorption experiments performed with 2ZC-cal without competitors (Sb), and with coexistent HCO3 <sup>−</sup> (Sb/HCO3), SO4 <sup>2</sup><sup>−</sup> (Sb/SO4) or HAsO4 <sup>2</sup><sup>−</sup> (Sb/HAsO4). (A<sup>n</sup><sup>−</sup> = HCO3 <sup>−</sup>, SO4 <sup>2</sup><sup>−</sup> or HAsO4 <sup>2</sup>−; na = not analyzed; dl = detection limit; dlZn = 0.2 μmol/L; dlAl = 7 μmol/L).

The effect of coexistent anions on Sb(OH)6 <sup>−</sup> uptake resulted to be HAsO4 <sup>2</sup><sup>−</sup> >> HCO3 <sup>−</sup> ≥ SO4 <sup>2</sup><sup>−</sup> for 2ZC-cal and, in partial agreement with previous work [30], HAsO4 <sup>2</sup><sup>−</sup> >> HCO3 <sup>−</sup> >> SO4 <sup>2</sup><sup>−</sup> for 3HT-cal.

At the end of the experiments with 3HT-cal, slight concentrations of Mg were determined, whereas the concentrations of Al and Fe were always below the corresponding detection limits (Table 3).

**Table 3.** Solution pH values and dissolved ions determined before and at the end of the sorption experiments performed with 3HT-cal without competitors (Sb), and with coexistent HCO3 <sup>−</sup> (Sb/HCO3), SO4 <sup>2</sup><sup>−</sup> (Sb/SO4) or HAsO4 <sup>2</sup><sup>−</sup> (Sb/HAsO4). (A<sup>n</sup><sup>−</sup> = HCO3 <sup>−</sup>, SO4 <sup>2</sup><sup>−</sup> or HAsO4 <sup>2</sup>−; na = not analyzed; dl = detection limit; dlAl = 7 μmol/L; dlFe = 1 μmol/L).


#### 3.1.2. Kinetics

In the experiments with both 2ZC-cal and 3HT-cal the solution pH values increased sharply (up to about 11) after the addition of sorbents (Figure 2a,b), and decreased after 48 h in the range of 8.1–9.8 (Tables 2 and 3; Figure 2a,b). Most of the Sb(OH)6 − was removed within the first six hours (Figure 2c,d) indicating that Sb(OH)6 − uptake occurred mainly during the reconstruction of the lamellar structure of LDHs [36] as schematized in reaction (1):

$$\begin{array}{c} \text{M}\_{1-x}^{2+} \text{M}\_{x}^{3+} \text{(OH)}\_{2} \text{ (CO}\_{3}\text{)}\_{\frac{x}{2}} \xrightarrow{\text{calcination}} \text{M}\_{1-x}^{2+} \text{M}\_{x}^{3+} \text{O}\_{1+\left(\frac{x}{2}\right)} + \text{xSb(OH)}\_{6}^{-} \\ \xrightarrow{\text{reconstruction}} \text{M}\_{1-x}^{2+} \text{M}\_{x}^{3+} \text{(OH)}\_{2} \text{(Sb(OH)}\_{6}\text{)}\_{x} + \text{xOH}^{-} \end{array} \tag{1}$$

The Sb(OH)6 − removal as a function of time was studied through the pseudo-first order [38] and the pseudo-second order kinetic models [39].

The Sb(OH)6 − sorption capacity has been calculated through the Formula (2):

$$\mathbf{q}\_t = (\mathbf{C}\_0 - \mathbf{C}\_t) \cdot \mathbf{V} / \mathbf{W} \tag{2}$$

where the sorption capacity (qt) is the amount of Sb(OH)6 − sorbed per unit of sorbent (mmol/g) at the reaction time t (h), C0 and Ct are the Sb(OH)6 − concentrations in solution (mmol/L) before the addition of the sorbent and at the reaction time t, V is the volume of solution (L) and W the weight of sorbent (g).

The pseudo-first order and the pseudo-second order equations are expressed as follows:

$$1\text{ pseudo} - \text{first order kinetic model} \qquad \qquad k\_1 = \frac{2.303}{t} \cdot \log \frac{q\_\varepsilon}{q\_\varepsilon - q\_t} \tag{3}$$

$$1\text{ pseudo}-\text{second order kinetic model}\qquad\qquad k\_2 = \frac{1}{t} \cdot \frac{q\_t}{q\_c(q\_c - q\_t)}\tag{4}$$

where qe is the Sb(OH)6 <sup>−</sup> sorption capacity at equilibrium (mmol/g), *k*<sup>1</sup> (1/h) and *k*<sup>2</sup> (g/mmol h) are the rate constant of sorption. These equations expressed in the linear form result as follows:

$$1\text{ pseudo} - \text{first order kinetic model}\qquad\qquad\qquad\frac{dq\_t}{dt} = k\_1 \left(q\_t - q\_t\right)\tag{5}$$

**Figure 2.** The solution pH values determined as a function of time during the sorption experiments without competitors and with coexistent anions performed with (**a**) 2ZC-cal and (**b**) 3HT-cal. The Sb(OH)6 − sorption capacity (qt) as a function of time in the experiments performed with (**c**) 2ZC-cal and (**d**) 3HT-cal without competitors and with coexistent anions.

To verify the applicability of kinetic models at the sorption system, the experimental data were plotted as log(qe − qt) vs. time for the pseudo-first order (Supplementary Material Figure S2) and t/qt vs. time (Supplementary Material Figure S3) for the pseudo-second order kinetic model. If the plots give a linear correlation, then the theoretical sorption capacity at equilibrium and the rate constants can be calculated from the slope and the intercept of the straight lines. The good fit of the data, the r<sup>2</sup> values close to the unit and the good agreement between the experimental sorption capacity at equilibrium (qe) and the theoretical sorption capacity (qcalc) indicated that the sorption system is better described by the pseudo-second order kinetic model (Table 4), suggesting that the Sb(OH)6 − uptake by both 2ZC-cal and 3HT-cal might principally occur by chemisorption.


**Table 4.** The Sb(OH)6 − sorption capacity of 2ZC-cal and 3HT-cal determined at equilibrium from the experimental data (qe) and calculated from the kinetic models (qcalc).

#### 3.1.3. Characterization of Sorbents

The XRD pattern of 2ZC-CO3 (Figure 3a) showed the characteristic basal reflections (003) and (006) attributable to a zaccagnaite-like compound [40].

**Figure 3.** XRD patterns of (**a**) 2ZC-CO3 and its calcined product 2ZC-cal and of (**b**) 3HT-CO3 and its calcined product 3HT-cal, and XRD patterns of sorbents recovered after the sorption experiment without competitors (Sb), and with coexistent SO4 <sup>2</sup><sup>−</sup> (Sb/SO4), HCO3 <sup>−</sup> (Sb/HCO3) or HAsO4 <sup>2</sup><sup>−</sup> (Sb/HAsO4). The blue dashed lines indicate the peaks of the (**a**) zincalstibite-like and (**b**) brandholzite-like compounds formed after the sorption experiments.

After calcination, in the XRD pattern of 2ZC-cal, the absence of LDH basal reflections and the presence of broad peaks at 32.1◦ and 36.4◦ 2θ, ascribable to ZnO, indicated the collapse of the lamellar LDH structure and the formation of a disordered ZnO, with Al probably dispersed in its structure [41]. Peaks of undesired phases were not detected. The XRD patterns of solids recovered after the sorption experiments showed two different LDH phases: the peaks at angular position 11.6◦ and 23.5◦ 2θ corresponded to the basal reflections (003) and (006) of 2ZC-CO3; instead, the peaks at 5◦ and 18◦ 2θ were attributable, respectively, to the (001) and (002) reflections of a zincalstibite-like compound [36,42].

The characteristic hydrotalcite-like compound basal reflections (003) and (006), visible in the XRD patterns of 3HT-CO3 (Figure 3b), were no longer detectable in the calcined phase (3HT-cal) which showed peaks at about 42◦ and 62◦ 2θ ascribable to a disordered MgO with the trivalent metals Fe and Al probably dispersed in its structure [26]. After the sorption experiments, the solids showed the characteristic peak of a brandholzite-like compound at 19.2◦ 2θ and a further brandholzite peak at 33.6◦ 2θ [36,43], except for the experiment with coexistent HAsO4 <sup>2</sup><sup>−</sup> where the characteristic brandholzite peak was barely visible and additional peaks at low angle compatible with the basal reflections of 3HT-CO3 were clearly recognizable.

*3.2. Sorption Experiments with Su Suergiu Mine Drainage (SU1)*

#### 3.2.1. Solutions

For convenience in this section the concentrations of ions in solution will be expressed as mg/L and μg/L.

The results of the chemical analysis of SU1 slag drainage showed a Ca-SO4 dominant chemical composition and high concentrations of Sb (9900 μg/L) and arsenic (As = 3390 μg/L) (Table 5).



The high value of EC (electrical conductivity) is related to the high contents of Ca2<sup>+</sup> and SO4 2−, with SO4 <sup>2</sup><sup>−</sup> mainly deriving from the oxidation of sulfides; moreover, the high concentration of HCO3 − avoids the decrease of pH, which resulted in slightly alkaline values (Table 5). The high concentration of both Sb and As is the consequence of the water interaction with the foundry slags [6,37]. It has been reported that in the Su Suergiu mine water, the Sb(III), when detected, results <2% of total dissolved Sb (water fraction < 0.45 μm) and that all the Sb(V) occurs as Sb(OH)6 − [6,44]. The concentration of Sb(III) determined in the SU1 sampled for sorption experiments with 3HT-cal and 2ZC-cal resulted to be 147 μg/L, therefore, in the present work, the Sb(OH)6 − is considered the only Sb form involved in the Sb removal processes. Excluding the experiments performed with 0.1 g of sorbents, at the end of reaction time up to 90–99% of Sb was removed from the solution, with 2ZC-cal slightly more effective than 3HT-cal (Figure 4a), whereas As was effectively removed in all experiments (Figure 4b). In almost all experiments, the Sb and As concentrations decreased close to, or below, the limits established for drinking water (Figure 4c,d).

**Figure 4.** Results of batch sorption experiments performed with 3HT-cal and 2ZC-cal with the Su Suergiu mine drainage (SU1). Percentages of (**a**) Sb and (**b**) As removed at the end of the sorption experiments, and concentrations of residual (**c**) Sb and (**d**) As in solution at the end of the experiments for different amounts of the sorbent. In the plots (**c**) and (**d**) the blue lines indicate the starting Sb or As concentrations; the red lines indicate the limits of Sb and As set for drinking water by the World Health Organization (WHO) [2] and the European Community (EU) [3].

#### 3.2.2. Sorbents

The XRD patterns of solids recovered after all experiments showed peaks at low angles compatible with the carbonate bearing LDHs, indicating the reconstruction of the typical lamellar LDH structure (Figures 5 and 6). In the range of pH values of the experiments, Sb and As prevail, respectively, as Sb(OH)6 <sup>−</sup> and HAsO4 <sup>2</sup>−, but peaks attributable to a brandholzite-like compound were not visible in the solids recovered after the experiments performed with 3HT-cal (Figure 5), and only after the experiment with 0.1 g of 2ZC-cal the characteristic peak of a zincalstibite-like compound was clearly recognizable (Figure 6). Therefore, as a consequence of the complexity of the chemical composition of SU1, the Sb(OH)6 <sup>−</sup> removal did not occur through the formation of Sb(OH)6 − bearing phases, but rather, it is reasonable to suppose that Sb(OH)6 − was incorporated in the interlayer region together with other anions, i.e., CO3 <sup>2</sup><sup>−</sup> and HAsO4 <sup>2</sup>−. It is also noticeable that all samples contained additional well defined peaks at about 30◦ 2θ ascribable to calcite (Figures 5 and 6), and after the experiment with 0.1 g of 2ZC-cal, the peaks attributable to monohydrocalcite were also present (Figure 6).

**Figure 5.** XRD patterns of 3HT-cal recovered after the sorption experiments with the slag drainage SU1 performed with different amounts of 3HT-cal (0.1, 0.25, 0.5, 1 g). On the upper right side of each XRD pattern the pH of solution as a function of time during each experiment is reported. The blue dashed lines indicate the basal reflection of layered double hydroxides (LDHs) formed after the sorption experiments.

**Figure 6.** XRD patterns of 2ZC-cal recovered after the sorption experiments with the slag drainage SU1 performed with different amounts of 2ZC-cal (0.1, 0.25, 0.5, 1 g). On the upper right side of each XRD pattern the pH of solution as a function of time during each experiment is reported. The blue dashed lines indicate the basal reflection of LDHs formed after the sorption experiments.

#### **4. Discussion**

#### *4.1. E*ff*ect of Coexistent Anions*

The results of mineralogical characterizations and chemical analysis, in agreement with previous works, showed that 2ZC-cal removed Sb(OH)6 − from solution through the reconstruction of a zincalstibite-like compound [36], and the 3HT-cal removed Sb(OH)6 − through the formation of a low ordered brandholzite-like compound [30,36]. The zincalstibite is an Sb(OH)6 − bearing LDH (with general formula Zn2Al(OH)6[Sb(OH)6] [25,42], whereas the brandholzite is a non-LDH phase with general formula Mg[Sb(OH)6]2·6H2O, whose the layered structure is characterized by the presence of two layers, {[Sb(OH)6]9} <sup>9</sup><sup>−</sup> and {[Sb(OH)6]3[Mg(H2O)6]6} <sup>9</sup>+, alternatively stacked along the *c* axis [43].

In agreement with the results of chemical analysis, the diffraction peaks of the Sb(OH)6 − bearing phases were more intense and well defined in the XRD patterns of solids recovered after the experiments wherein the greatest amounts of Sb(OH)6 − were removed from solution. In the experiments performed with 3HT-cal and HCO3 <sup>−</sup> and HAsO4 <sup>2</sup><sup>−</sup> as coexistent anions, the decrease in Sb(OH)6 − removal was linked to the appearance of the (003) and (006) basal reflections of LDHs (Figure 3b). In particular, in the experiment with coexistent HCO3 − the competition effect was improved by the increase of solution pH values, up to 10.6 (Figure 2b), that favors the prevalence of CO3 <sup>2</sup><sup>−</sup> in solution that has a high affinity for the LDH interlayer [23,30]. Also HAsO4 <sup>2</sup><sup>−</sup> has a high affinity for hydrotalcite-like compounds [27,45], but previous authors have observed that the removal capacity of calcined MgAl-LDHs is higher for Sb(OH)6 <sup>−</sup> than for HAsO4 <sup>2</sup>−, suggesting that the uptake of Sb(OH)6 − through the selective

crystallization of a brandholzite-like compound is more favorable than the sequestration from solution by the intercalation in the interlayer [46]. The results of our work showed that, when coexisting in solution, HAsO4 <sup>2</sup><sup>−</sup> strongly competes with Sb(OH)6 − and is preferentially removed, probably due to its higher specific ionic charge.

The low amounts of Mg determined at the end of experiments with 3HT-cal indicated a slight dissolution of sorbent (6–15%) (Table 3). Because the concentrations of dissolved Fe and Al were always below the corresponding detection limits and no Al and/or Fe secondary phases were observed in the XRD patterns of solids recovered at the end of experiments (Figure 3b), the low amount of Al and Fe released in the solution might precipitate as amorphous solids. The Fe and Al of undissolved phase that removed the Sb(OH)6 − through the formation of brandholzite-like compound, probably remained in undetermined sites of the low ordered brandholzite structure.

The XRD pattern of the experiment performed with 2ZC-cal and HAsO4 <sup>2</sup><sup>−</sup> as coexistent anion showed well defined basal reflections compatible with the original 2ZC-CO3 LDH, while the characteristic zincalstibite-like compound peaks were scarcely visible (Figure 3a). Moreover, two broad undefined humps in the angular ranges 30–33◦ and 35–37◦ 2θ indicated that, at the end of the experiment, part of the 2ZC-cal did not react and explained the low Sb(OH)6 − removal that, unlike the experiment performed with 3HT-cal, cannot be attributable to the preferential uptake of HAsO4 2− (Table 2). Previous authors reported the high affinity of As(V) for sulfate bearing ZnAl-LDHs [47,48] but, at the best of our knowledge, experiments on the As(V) removal by calcined ZnAl-LDHs is lacking and needs further study.

#### *4.2. Sorption Experiments with Su Suergiu Mine Drainage (SU1)*

The results suggested that both 2ZC-cal and 3HT-cal are suitable for the Sb (and also As) removal from SU1; however, to assess their practical use, the overall quality of treated water must also be considered. The decrease of EC value after the experiments can be attributable to the decrease of Ca2<sup>+</sup> and HCO3 <sup>−</sup>. The contents of Na+, K<sup>+</sup> and Cl<sup>−</sup> did not vary significantly, instead sensible variations of SO4 <sup>2</sup><sup>−</sup>, NO3 − and F− were observed in a few cases. The partial dissolution of sorbents explained the irregular increase of dissolved Zn or Mg and Al. The limit of Al in drinking water is established at 0.2 mg/L by both the WHO and EU, whereas the concentration of Zn in drinking water is not regulated by the WHO and EU, but rather the Italian Legislation set 2 mg/L [49]; therefore, the concentration of Al exceeded the limit only in one experiment and the Zn limit was never reached (Table 5).

In the experiment performed with 0.1 g of 3HT-cal the amount of Sb removed was dramatically lower with respect to the other experiments with higher amounts of 3HT-cal. In this case it was possible to observe that the Sb removal did not increase with the weight of 3HT-cal used. The Ca2<sup>+</sup> and HCO3 <sup>−</sup> decreases did not show correlation (Supplementary material Figure S4). In particular, the Ca2<sup>+</sup> concentration suggested a lower CaCO3 precipitation at the end of experiments with 0.1 and 1 g of 3HT-cal with respect to the other ones. In the first case (i.e., 0.1 g of 3HT-cal) the CaCO3 precipitation should be limited by the low increase of pH (8.2–8.4), while in the experiment with 1 g of 3HT-cal, where the pH values increase up to 10.3 (Figure 5) and CO3 <sup>2</sup><sup>−</sup> prevails in solution, the CaCO3 precipitation should be hindered by the uptake of CO3 <sup>2</sup><sup>−</sup> in the interlayer during the 3HT-cal rehydration (Figure 5). This could also explain the low Sb removal in spite of the high amount of 3HT-cal. In fact, at high pH the Sb uptake was hindered by CO3 <sup>2</sup><sup>−</sup> that strongly competes for the entry in the interlayer region of the reconstructing lamellar LDH structure. The experiments with 0.5 and 0.25 g of 3HT-cal seemed the best compromise to reach the most favorable conditions for the highest Sb and As removal from 400 mL of SU1.

As observed above, at the end of the experiments with 0.25, 0.5 and 1 g of 2ZC-cal, the amount of Sb removed from solution was markedly higher with respect to 0.1 g of sorbent. Moreover, it is possible to observe that the dissolved concentrations of Sb and As slightly decreased as the amount of 2ZC-cal used increased. The decrease of Ca2<sup>+</sup> and HCO3 − observed at the end of the experiments did not show clear correlation (Supplementary material Figure S4). The sequestration of Ca2<sup>+</sup> is attributable to the precipitation of calcium carbonates, which can also explain the slight decrease of Mg2<sup>+</sup> in some cases (Table 5). The slight differences in residual Ca2<sup>+</sup> concentration indicated the precipitation of nearly equivalent amounts of calcium carbonates in the different experiments (Table 5). Differently, the HCO3 − concentration decreased as the amount of 2ZC-cal increased because its uptake from the solution occurred by both the precipitation of calcium carbonates and the entry in the interlayer during the LDHs reconstruction (Supplementary material Figure S4). It is possible to note that only in the experiments with 0.5 and 1 g of 2ZC-cal the pH of solution slightly increased after the addition of sorbents (up to 8.7) and successively decreased to 7.9–8.0, whereas in the experiments with 0.1 and 0.25 g, the variation of pH values was negligible and remained in the range of 8.1–8.4 (Figure 6). These limited pH variations among the experiments with 2ZC-cal can explain the precipitation of nearly equivalent amounts of calcium carbonates. The relatively low pH values can also explain the limited precipitation of calcium carbonates, deducible from the concentrations of Ca removed from solution (Supplementary material Figure S4), with respect to that observed in the experiments performed with 3HT-cal. At these pH values around 8, HCO3 − prevails among the dissolved carbonate species, therefore it is possible that the reconstruction of LDH by rehydration of 2ZC-cal occurred via the intercalation of HCO3 <sup>−</sup>, as well as CO3 <sup>2</sup>−.

In our previous work we performed sorption experiments, carried out with ultrapure water containing only Sb(OH)6 <sup>−</sup>, to determine the maximum theoretical Sb(OH)6 <sup>−</sup> sorption capacity (*qmax*) of 2ZC-cal and 3HT-cal through the Langmuir isotherm [36]. The values of *qmax* were 4.37 mmol/g (i.e., 532 mg/g) for 3HT-cal and 4.54 mmol/g (i.e., 553 mg/g) for 2ZC-cal [36]. In the present work we have observed that the coexistence of other anions in solution can affect the Sb(OH)6 − removal capacity of sorbents tested. In order to construct the isotherm, being the starting Sb concentration of sorption experiments that of SU1, the solid/liquid ratio has been changed. However, the sorption data did not fit for the calculation of the isotherm because, due to the complexity of the SU1 chemical composition, also other processes, like the precipitation of calcium carbonates, occurred during the interaction between water and calcined LDHs having an effect on the Sb removal.

It is worth mentioning that the LDHs exhibit the possibility of being reused by regeneration, operated through calcination or anion exchange, for consecutive sorption–regeneration–sorption cycles [50–54]. This is a very important characteristic for their practical use in water treatment because it can reduce the amount of post-treatment waste materials. In this regard, as far as we know, there is no data about the regeneration of ZnAl-LDH after Sb(V) adsorption; whereas it has been reported that, because brandholzite is a non-LDH mineral, the brandholzite formation connected with the Sb(V) removal by calcined MgAl-LDHs can negatively affect the MgAl-LDH regeneration capacity [46,55]. In this work we have observed that in the experiments performed with SU1, the Sb(V) removal by 3HT-cal occurred by intercalation in the LDH interlayer. It is reasonable to suppose that this removal mechanism may positively influence the effectiveness of LDHs regeneration. Therefore, further studies should be performed in order to assess the potential use of both 3HT-cal and 2ZC-cal regenerated after Sb(V) removal from real polluted water.

#### **5. Conclusions**

In this work the Sb(V) (in the Sb(OH)6 − form) removal capacity of calcined hydrotalcite-like (3HT-cal) and zaccagnaite-like (2ZC-cal) compounds has been studied in order to assess their potential for the practical use with real Sb(V) polluted water. For this purpose, first the effect of other anions on the Sb(OH)6 − removal capacity of 2ZC-cal and 3HT-cal were tested through batch sorption experiments with coexistent anions in solution. Successively batch sorption experiments were carried out with the slag drainage (SU1) sampled at the abandoned Su Suergiu mine (Sardinia, Italy) affected by relevant Sb pollution.

In agreement with previous studies, the results of our experiments with coexistent anions showed that 3HT-cal and 2ZC-cal removed Sb(OH)6 − through the formation of brandholzite-like and zincalstibite-like compounds, respectively. Among the anions tested, the competition effect on Sb(OH)6 <sup>−</sup> removal resulted to be HAsO4 <sup>2</sup><sup>−</sup> >> HCO3 <sup>−</sup> ≥ SO4 <sup>2</sup><sup>−</sup> for 2ZC-cal, and HAsO4 <sup>2</sup><sup>−</sup> >> HCO3 − >> SO4 <sup>2</sup><sup>−</sup> for 3HT-cal.

The results of the sorption experiments showed that both 3HT-cal and 2ZC-cal effectively removed Sb and As from SU1 and, thus, these phases might have potential use for practical application with real polluted water. In the case under study, due the high concentration of carbonate species in SU1, Sb was mainly removed by intercalation in the interlayer of carbonate bearing LDHs rather than by the formation of zincalstibite-like and brandholzite-like compounds, as instead observed in the competition experiments performed with synthetic solutions. The results also indicated that the precipitation of calcium carbonates (i.e., calcite and monohydrocalcite) may favor the Sb removal, subtracting CO3 2− from the solution as a possible strong competitor of Sb for the interlayer region of LDHs. Therefore, especially in the experiments with 3HT-cal, the Sb removal capacity was markedly influenced by the liquid/solid ratio that determines the solution pH, the correlated precipitation of calcium carbonates and the competition effect of carbonate species in solution. The Sb removal by intercalation in the interlayer of carbonate bearing LDHs rather than the formation of antimonate bearing phases may represent an advantage for the LDHs regeneration, therefore further studies should be addressed in that direction.

**Supplementary Materials:** The following are available online at http://www.mdpi.com/2073-4352/9/8/410/s1, Figure S1: The abandoned mine area of Su Suergiu (Cidu et al [34], modified); Figure S2: The linear kinetic plots of the pseudo-first order rate equation of Sb(OH)6 − sorption experiments performed without competitors and with coexistent anions performed with (**a**) 2ZC-cal and (**b**) 3HT-cal; Figure S3: The linear kinetic plots of the pseudo-second order rate equation of Sb(OH)6 − sorption experiments without competitors and with coexistent SO4 <sup>2</sup><sup>−</sup> and HCO3 <sup>−</sup> performed with (**a**) 2ZC-cal and (**b**) 3HT-cal, and (**c**) performed with HAsO4 <sup>2</sup>−; Figure S4: (**a**) Plot of Ca removed vs HCO3 removed, and concentration of (**b**) Ca and (**c**) HCO3 removed at the end of the experiments performed with 3HT-cal or 2ZC-cal and SU1. The horizontal green lines in the plots (**b**) and (**c**) indicate the starting concentrations of Ca or HCO3, respectively.

**Author Contributions:** Conceptualization, E.D., F.F. and R.C.; Data curation, E.D.; Formal analysis, E.D.; Funding acquisition, F.F. and R.C.; Investigation, E.D.; Methodology, E.D.; Supervision, F.F. and R.C.; Validation, F.F. and R.C.; Writing—original draft, E.D.; Writing—review and editing, E.D., F.F. and R.C.

**Funding:** Authors thank the financial support from the Ministero Università Ricerca Scientifica Tecnologica through the research projects PRIN 2009 (Coordinator R. Cidu) and PRIN 2010–2011 (Coordinator P. Lattanzi).

**Acknowledgments:** Authors thank the Paola Meloni and Ombretta Cocco of "Research and Didactic Laboratory for the Conservation of the artistic heritage of the Monumental Cemetery of Bonaria" (Department of Mechanical, Chemical and Materials Engineering (DIMCM), University of Cagliari) for the assistance with the XRD measurements.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


© 2019 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (http://creativecommons.org/licenses/by/4.0/).

## *Article* **Mg-Fe Layered Double Hydroxides Enhance Surfactin Production in Bacterial Cells**

**Pei-Hsin Chang 1, Si-Yu Li 1,2, Tzong-Yuan Juang 3,\* and Yung-Chuan Liu 1,\***


Received: 20 June 2019; Accepted: 11 July 2019; Published: 12 July 2019

**Abstract:** In this study, four additives—montmorillonite, activated carbon, and the layered double hydroxides (LDHs), Mg2Fe–LDH and Mg2Al–LDH—were tested for their ability to promote surfactin production in a *Bacillus subtilis* ATCC 21332 culture. Among these tested materials, the addition of 4 g/L of the Mg-Fe LDH, which featured an Mg/Fe molar ratio of 2:1, produced the highest surfactin yield of 5280 mg/L. During the time course of *B. subtilis* cultivation with the added LDH, two phases of cell growth were evident: Growth and decay. In the growth phase, the cells grew slowly and secreted a high amount of surfactin; in the decay phase, the cells degraded rapidly. The production in the presence of the Mg2Fe–LDH had three characteristics: (i) High surfactin production at low biomass, indicating a high specific surfactin yield of 3.19 g/g DCW; (ii) rapid surfactin production within 24 h, inferring remarkably high productivity (4660 mg/L/d); and (iii) a lower carbon source flux to biomass, suggesting an efficient carbon flux to surfactin, giving a high carbon yield of 52.8%. The addition of Mg2Fe–LDH is an effective means of enhancing surfactin production, with many potential applications and future industrial scale-up.

**Keywords:** *Bacillus subtilis*; surfactin; quantitative analysis; fermentation; growth phase; layered double hydroxides

#### **1. Introduction**

Biosurfactants are amphipathic molecules produced by microorganisms [1–3] with the capability of decreasing surface and interfacial tension [4]. Depending on their chemical composition and their producing organism, biosurfactants can possess high biodegradability, low toxicity, ecological acceptability, and high efficiency. Accordingly, they have been investigated as possible alternatives to chemical surfactants [5,6]. *Bacillus* spp., bacterial strains of complicated physiological diversity, can be used to produce many bioactive peptides with potential biotechnological and biopharmaceutical applications. Among these peptides, the lipopeptides that feature an alkyl group and a circular peptide group are the most popular biosurfactants [7]; these materials include surfactins [8–10], iturins [11,12], and fengycins [13].

The surfactin produced by *B. subtilis* is one of the strongest biosurfactants available [7]. Its chemical composition is that of a cyclic lipopeptide (comprising seven amino acids) with a 12 to 19-carbon atom hydrophobic fatty acid chain [14]. Surfactin can lower the surface tension of water to 27 mN/m even when its concentration is as low as 0.005% [7,10,15,16], suggesting its great potential applicability. Nevertheless, the high expense and low yield of surfactin production have limited its commercial use. Yeh et al. found that limiting the concentration of the carbon source (glucose) affected the surfactin production mediated by *B. subtilis* [17]. Davis et al. observed the highest production of surfactin when ammonium nitrate was the nitrogen source during *B. subtilis* cultivation in a defined medium [18]. Sen et al. noted that the ratio of Mn and Fe mineral salts in the medium was a factor affecting the production of surfactin [19]. Wei and Chu found that the yield of surfactin increased dramatically, over those obtained using genetic strains, when employing 0.01 mM Mn2<sup>+</sup> [20]. Furthermore, Wei et al. employed an iron-enriched (4 mM Fe2<sup>+</sup>) minimal salt medium to produce 3000 mg/L of surfactin [21]. Moreover, some of these studies revealed that the addition of solid additives (e.g., activated carbon (AC) or expanded clay) could increase surfactin production significantly. For example, Yeh et al. added AC and increased the yield of surfactin to 3600 mg/L [17].

Layered double hydroxides (LDHs)—also known as anionic clays—comprise cationic brucite-like layers with exchangeable interlayer anions [22]. Because a positive ionic charge appears on the surface layer, many types of molecules can be intercalated into LDHs [23–27]. Several methods have been developed to widen the layered gallery, with globular macromolecules as intercalating agents [28,29]. Conterosito et al. intercalated various pharmaceutics drugs and cosmetic sunscreen into Mg-Al\_LDH and Zn-Al\_LDH. They revealed that different bioactive molecules could interact with inorganic LDH and demonstrated the relationship between the molecular length and an enlarged interlayer spacing [30]. Toson et al. showed the intercalation of organic molecules into the LDH interlayer by the liquid-assisted grinding method. The intercalation mechanism for layer widening with intercalated organic molecules was investigated [31]. Choy et al. employed supramolecular inorganic species (e.g., nanoscale Mg-Al LDH) as biomolecule reservoirs that could be used for gene and drug delivery [23,32]. Nevertheless, few studies have focused on using LDHs as additives for microbial cultivation. Kan et al. prepared Mg2Al–LDH and investigated its effect as an additive on surfactin production and surfactin intercalation [33]. When considering the application of the surfactin-intercalated LDH as a slow release bio-pesticide, however, this aluminum salt was prohibited from field tests [34]. For agricultural applications, iron salts are generally considered less toxic. Therefore, in this present study, we prepared several Mg-Fe LDH derivatives with potentially greater practicality. We tested the effects of their addition on the production of surfactin from a *B. subtilis* culture. To our surprise, replacing the additive to Mg2Fe–LDH had an extraordinary effect on the surfactin production. Accordingly, we examined various MgnFe–LDH (n = 1, 2, 3) compositions and concentrations to determine the optimal conditions for surfactin production. In addition, we examined the time course of the production in the optimal culture. An extra low biomass of cells yielded the highest surfactin production. This result was quite different from that obtained after the addition of Mg2Al–LDH. Furthermore, we compared the effects of the LDHs with those of other additives (e.g., montmorillonite (MMT), AC), and determined the conditions for the highest production of surfactin through quantitative analysis. Herein, we also suggest possible reasons for the enhancement of surfactin production mediated by LDHs.

#### **2. Materials and Methods**

#### *2.1. Chemicals and Reagents*

Al(NO3)3·9H2O, Fe(NO3)3·9H2O, and Mg(NO3)2·6H2O were purchased from SHOWA, USA. MMT was obtained from Alfa Aesar, USA. The AC was obtained from China Activated Carbon (Taipei, Taiwan); it had a diameter of 3 to 4 mm, a height of 9 mm, and a specific surface area of 1200 m2/g, and was prepared from bituminous coal with an iodine number of 1150 mg/g. Surfactin (≥98%, Sigma–Aldrich, Missouri, MO, USA) was used as the standard. All solvents and other chemicals were of analytical grade.

#### *2.2. Microorganisms and Culture Conditions*

The strain, *B. subtilis* ATCC 21332, was obtained from Professor Wei Yu-Hong of Yuan Ze University. This strain was kept on a nutrient-agar plate at 30 ◦C. For cultivation, its seed medium comprised 1% glucose, 0.5% yeast extract, 1% peptone, and 1% NaCl. The seed culture was performed in Erlenmeyer flasks (500 mL) containing the seed medium (100 mL) inoculated with two loops of cells. The cultivation was conducted at 200 rpm and 30 ◦C for 12 h. The main shake-flask culture was

conducted in an Erlenmeyer flask (500 mL) containing the main medium (100 mL) comprising 10 g/L sucrose, 5 g/L (NH4)2SO4, 5.67 g/L Na2HPO4, 4.08 g/L KH2PO4, 0.2 g/L MgSO4·7H2O, and 0.57 g/L FeSO4·7H2O. The media were sterilized (121 ◦C, 20 min); the carbon source was autoclaved separately. The medium (90 mL) was inoculated with the seed broth (10 mL). The flasks were incubated on a rotary shaker (200 rpm, 30 ◦C, 5 days). When testing additives, MMT, AC, and the prepared LDHs were added (2 g/L) to the culture medium at the beginning of the culture process.

#### *2.3. Mg2Al–LDH and Mg2Fe–LDH*

Mg2Al-NO3–LDH and Mg2Fe-NO3–LDH were prepared through co-precipitation, as described previously [27]. To prepare the Mg2Al–LDH sample, Mg(NO3)2·6H2O (120 g) and Al(NO3)3·9H2O (90 g) were dissolved in deionized H2O (1 L). To prepare the Mg2Fe–LDH sample, Mg(NO3)2·6H2O (169 g) and Fe(NO3)3·9H2O (134 g) were mixed in deionized H2O (1 L). To prepare samples with Mg/Fe molar ratios of 1.0 and 3.0, appropriate amounts of Mg(NO3)2·6H2O and Fe(NO3)3·9H2O were used. Each aqueous solution was stirred vigorously at 80 ◦C while purging with nitrogen gas. When preparing the Mg2Al–LDH sample, the pH was maintained at 10 ± 0.2 by adding 4 N NaOH in portions. For the Mg2Fe-NO3–LDH sample, the pH was adjusted to 9.5 ± 0.2 by using a mixture of NaOH and K3[Fe(CN)6], prepared based on the following compositions: [OH–]/([Mg2+] + [Fe3+]) = 1.6 and [[Fe(CN)6] 3–]/[Fe3<sup>+</sup>] = 3. The suspension that formed was stirred at 80 ◦C for 24 h. The obtained precipitates—white Mg2Al–LDH and dark-red Mg2Fe–LDH—were filtered off and washed (deionized H2O). The filtered cakes were lyophilized (freeze-drying). The dried LDHs were characterized using x-ray diffraction (XRD; PANalytical, X'Pert PRO MRD, Almelo, Netherlands) and attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopy (Thermo Scientific, Nicolet iS50 FTIR, Madison, WI, USA).

#### *2.4. Quantitative Analysis*

To study the effects of LDH addition, the following quantitative terms are defined. The surfactin yield (mg/L) is expressed by the volumetric concentration. The carbon source yield is defined as:

$$\mathcal{Y}\_{\frac{\beta}{S}} = \frac{\Delta P}{\Delta S} \tag{1}$$

The productivity is defined as:

$$\text{Productivity} = \frac{\Delta P}{\Delta t} \tag{2}$$

The specific yield is defined as:

$$\mathcal{Y}\_{\frac{P}{\Delta}} = \frac{\Delta P}{\Delta X} \tag{3}$$

where *P* represents the surfactin concentration, *S* is the carbon source concentration, *X* is the concentration of biomass dried cells, and *t* is the duration of cultivation. All concentrations are expressed herein on a volumetric basis.

#### *2.5. Assays*

The surfactin concentration was measured using a modified approach called salt-assisted homogeneous liquid-liquid extraction via high-performance liquid chromatography (HPLC) [33,35]. The culture broth (1 mL) was subjected to a centrifugation (3200 g, 3 min, 4 ◦C) to remove the solid pellets. The supernatant was mixed with MeCN (0.5 mL) and ammonium sulfate (0.8 g) and subjected to vigorous stirring for 1 min, and then centrifuged (3200 g, 3 min). The supernatant was filtered (0.22 μm) to obtain the sample for injection. HPLC analysis was performed under the following conditions: A reversed-phase C-18A column (5 mm, 18 mm × 100 mm BDS-Hypersil, Thermo Fisher Scientific, Waltham, MA, USA); a mobile phase comprising CF3CO2H, MeCN, and deionized H2O (0.1:400:100); a flow rate of 1.0 mL/min; an injection sample volume of 20 μL; and a UV–Vis detector

(JASCO, Tokyo, Japan) operated at 220 nm. A standard curve was constructed using a freshly prepared solution of surfactin (Sigma). The chromatogram of the standard (supplemental Figure S1) revealed various ratios of the surfactin isoforms A–F. The surfactin produced using *B. subtilis* ATCC 21332 featured the same surfactin isoforms A–F at various ratios. In the surfactin assay, the whole isoforms were measured and added up for quantitative calculation. To analyze the cells' dried weight (CDW), 5 mL of the broth sample was subjected to centrifugation (12,000 *g*, 10 min) to obtain a pellet. Distilled H2O (5 mL) was added to the pellet; after adjusting to pH 2.0, the sample was vigorously stirred (1 min). The mixture was centrifuged (12,000 *g*, 5 min). The pellet obtained was dissolved in distilled H2O (5 mL); the pH was adjusted to 7.0 and the mixture was again subjected to centrifugation. The obtained pellet was washed with distilled H2O (2 × 5 mL), dried (80 ◦C, 12 h), and then weighed. The basal spacing of the LDH was determined using a Shimadzu SD-D1 X-ray diffractometer with a Cu target (scanning rate: 1◦/min). The basal spacing was estimated using the Bragg equation (*n*λ = 2*d*sinθ).

#### *2.6. Statistical Analysis*

Multiple flasks were run concurrently. Three flasks were employed each time for daily sampling. Each data point is expressed as a mean plus standard deviation. The Tukey test was applied for the comparison of results (*p* ≤ 0.05).

#### **3. Results and Discussion**

#### *3.1. Preparation of MgFe–LDH*

The addition of a small quantity of a solid carrier (AC or expanded clay) has been claimed as an effective approach toward increasing surfactin production [17]. LDHs are layered anionic exchange substances that have been intercalated with various macromolecules for the purpose of their slow release [36,37]. The addition of Mg2Al–LDH LDH to a surfactin production fermentation system involving *B. subtilis* incubation revealed that surfactin could indeed intercalate into the LDH layer gallery to form a surfactin–LDH complex; this phenomenon occurred with a significant increase in the production of surfactin [33]. In consideration of a slow-release composite for agricultural use, Mg2Al–LDH would be inappropriate for field trials. For this study, therefore, we prepared Mg2Fe–LDH instead. We examined the effect of adding this iron salt LDH to *B. subtilis* cultivation to study whether it, too, would promote surfactin production. The prepared Mg2Fe–LDH was subjected to XRD and ATR-FTIR spectroscopic analysis. These analyses revealed an Mg2Fe–LDH layer spacing of 7.8 Å at a value of 2θ of 11.3◦, derived from the calculation of Bragg's equation (Figure 1a), and a typical adsorption peak (1381 cm<sup>−</sup>1) for NO3 – anions within the prepared LDH (data not shown). In addition, to confirm the interaction between LDH and bacterial cells, the LDH after the cultivation was collected and subjected to XRD analysis. The result in Figure 1b shows that the collected LDH did vary its 2θ from the original 11.3◦ to 8.3◦, indicating a d-spacing of 10.8 Å. The original XRD peak with a d-spacing of 7.8 Å completely disappeared. The enlarged spacing was likely due to the LDH interaction with surfactin molecules. The isoelectric point (IEP) of surfactin is around pH 5, and the fermentation process while applying LDH to the cultivation was around pH 7.4. The pH higher than the IEP would allow the surfactin to possess a negative charge, giving the chance of anion exchange for LDH intercalation. Besides, the interlayer spacing expansion of LDH might be ascribed not only to the surface interaction of surfactin intercalation but also the combination of water and other anion molecules in the culture medium into the Mg2Fe–LDH interlayer.

**Figure 1.** XRD pattern of (**a**) pristine Mg2Fe–LDH and (**b**) Mg2Fe–LDH collected after fermentation.

#### *3.2. E*ff*ect of Solid Additives on Surfactin Production*

As reported previously, the addition of some solid additives can enhance surfactin production [17]. For this present study, four solid additives—MMT, AC, and two LDHs—were prepared and added respectively to the *B. subtilis* culture medium; the medium prepared without any additives was used as the control during the five-day fermentation. The surfactin production increased when the culture medium contained each of these solid additives, relative to the control. The addition of MMT, AC, and the two LDHs (2 g/L) resulted in surfactin yields that had increased by 2.0-, 3.0-, 3.8-, and 4.5-fold, respectively, when compared with the control (Figure 2). It is noteworthy that the AC with the alkaline characterization might lead to surfactin linearization and surfactin binding on the AC surface, which may be an underestimation of the actual production. Thus, the LDHs were the most effective carriers for enhanced surfactin production in a culture of *B. subtilis* ATCC 21332. Furthermore, the amount of surfactin produced in the presence of Mg2Fe–LDH was more than that produced in the presence of Mg2Al–LDH. Indeed, Mg2Fe–LDH had an extraordinary stimulatory effect on promoting surfactin production.

#### *3.3. E*ff*ect of MgFe–LDH Composition on Surfactin Production*

To study the effect of the Mg/Fe molar ratio on surfactin production, LDHs were prepared with Mg:Fe molar ratios of 1:1, 2:1, and 3:1 and added into *B. subtilis* cultivation. The concentrations of the additive ranged from 1 to 6 g/L in the fermentation medium. The cultivation was performed for 5 days. Figure 3 reveals that the LDH prepared with a Mg:Fe molar ratio of 2:1 had the greatest effect at promoting surfactin production. In general, LDHs possessing different ratios of divalent and trivalent metal ions possess different types of positively charged sheets and different layer dimensions in their resulting layered structures [38–41]. In the brucite-like layers of an LDH, a fraction of the divalent metal ions is replaced by trivalent metal ions, with the molar ratio of M3<sup>+</sup>:(M3<sup>+</sup> + M2+) (*x*) normally positioned between 0.2 and 0.4 [24,42]. In this present study, an Mg:Fe ratio of 2:1 (*x* = 0.33) had the best effect on improving surfactin production. Thus, it appears that the layer size associated with the positively charged sheets of the Mg2Fe–LDH structure had the strongest stimulatory effect on the cells.

**Figure 2.** Effects of solid additives on surfactin yield in *B. subtilis* ATCC 21332 cultivation in a 5-day fermentation: (**A**) none; (**B**) MMT; (**C**) AC; (**D**) Mg2Al–LDH; (**E**) Mg2Fe–LDH. The surfactin level in the supernatant of the broth was determined. Error bars indicate the standard deviations from three tests.

**Figure 3.** Effect of Mg2Fe–LDH concentration (1, 2, 4, and 6 g/L) on the specific surfactin yield, where bars with different patterns represent various Mg2<sup>+</sup>:Fe3<sup>+</sup> ratios: blank, 1:1; back slash, 2:1; slash, 3:1. Error bars indicate standard deviations from three tests.

#### *3.4. Time Course of Cultivation with LDH Addition*

To study the cell growth after adding LDH, the time courses of the cultivation events performed with and without added LDH were recorded (Figure 4). Although the addition of Mg2Fe–LDH promoted surfactin production, relative to that of the control, it was interesting to observe that the cell growth ended on the first day, where the amount of surfactin reached 4.8 g/L. In terms of product formation kinetics, this behavior was a clear growth-associated pattern: The cells grew and surfactin was produced. After day 1, the cells began to degrade in a decay phase, with the surfactin production decelerating. In contrast, the growth of cells was very rapid in the culture medium prepared without LDH, but the level of surfactin production was very low. Thus, a slight inhibition of cell growth appeared to trigger the cells to secrete more surfactin. We suspect that the surfactin secreted by the cells performed a role as a protecting agent that kept the cells from coming into direct contact with the LDH.

**Figure 4.** Time courses of surfactin and biomass production in the presence and absence of Mg2Fe–LDH (4 g/L). Error bars indicate the standard deviations from three tests.

Figure 5A,B present microscopy images of the morphologies of the cells grown in the presence of the LDH. In the culture medium lacking the LDH, the cells had a short and rod-like morphology from day 1 to day 3 of culturing. By the fifth day, some cells became slenderer than the original short-rod cells. In contrast, in the culture medium incorporating the LDH, the cells grew in a short-rod shape on the first day, but, by the third day, most of the cells had decayed and shrunk, with many endospores present. By the fifth day, almost none of the cells were evident in the broth, with only some spores remaining in the culture. This observation is consistent with the amounts of cells measured in the study. Therefore, the addition of LDH did inhibit the growth of cells during the cell growth phase, but it also enhanced the production of surfactin. Accordingly, in addition to the high surfactin yield of the culture incorporating the LDH, an extremely high productivity also ensued. Because of the lower number of cells, not only was the specific production elevated, the carbon source conversion to surfactin was also enhanced and provided a high carbon yield.

**Figure 5.** Optical microscopy images of *B. subtilis* ATCC 21332 growth, taken after various numbers of cultivation days; (**A**) in the absence of any LDH; (**B**) in the presence of Mg2Fe–LDH.

#### *3.5. Comparison of Surfactin Production*

Table 1 compares the surfactin production in this present study with those reported previously in the literature. Four factors characterize surfactin production in these bioprocesses in terms of their efficiency for fermentation on industrial scale: the surfactin yield, the carbon yield, the productivity, and the specific production. Due to the variation on surfactin quantification, the surfactin assays, such as HPLC, surface tensions, and acid precipitation, were also listed. As evident in Table 1, the addition of Mg2Fe–LDH had a unique effect in promoting surfactin production. Historically, surfactin production has improved gradually from an original yield of less than 1000 mg/L two decades ago to approximately 2000 to 3000 mg/L recently. When using this present approach, the yield of surfactin after the addition of Mg2Fe–LDH was enhanced significantly, to greater than 5000 mg/L. Furthermore, the addition of Mg2Fe–LDH ensured that the carbon source mostly flowed to surfactin production. Indeed, the carbon source yield was approximately 52.8%. This high carbon yield characterizes a surfactin production process with a highly efficient use of the raw material. In addition, the presence of Mg2Fe–LDH caused the surfactin yield to reach 4660 mg/L after one day of culturing; that is, the productivity was 4660 mg/L, a remarkably high value as compared in the literature. In addition, because the number of cells decreased in the presence of Mg2Fe–LDH, the smaller amount of biomass and the higher surfactin yield led to a specific yield of 3.19 g/g DCW. The addition of Mg2Fe–LDH in *B. subtilis* submerged cultivation provided a high carbon yield, high productivity, and high specific production of surfactin; such a high efficiency appears well suited to industrial applications.

At the beginning of our approach, the change of Mg2Al–LDH to Mg2Fe–LDH was due to the practical need in agricultural applications, where the aluminum salt is prohibited from field tests. However, to our surprise, the replacement of additive to Mg2Fe–LDH did give an extraordinarily high surfactin production. Due to this effect, the three critical characteristics affecting surfactin production were evaluated. It was found that a high specific surfactin yield, a high productivity, and a high carbon yield could be obtained in the presence of the Mg2Fe–LDH. To explain the difference between Mg2Fe–LDH and Mg2Al–LDH additions, the effect of the Mg2Fe–LDH addition with the leaking iron trace element in the culture was the possible reason for this highly efficient surfactin production. To decipher the cause of the extraordinarily high stimulatory effect of Mg2Fe–LDH, the following considerations might be taken into account. In some previous studies, ferric ions have been found to serve as trace element stimulators, with an excellent ability to promote surfactin production [21,43,44]. In addition, the use of pristine Mg2Al–LDH has been claimed to enhance surfactin production as a result of its toxicity toward the cells [33]. Accordingly, the presence of Mg2Fe–LDH was expected to not only inhibit cell growth and promote surfactin production (similar to the behavior of Mg2Al–LDH)

but also to slowly release some iron salts to serve as trace elements in the medium, thereby also improving the surfactin production. The higher production obtained using Mg2Fe–LDH, compared with that of Mg2Al–LDH, might be due to the synergistic effect of the Mg2Fe–LDH crystalline structure and the trace iron salts in the medium, with both combining to promote surfactin production to such a high level.


**Table 1.** Various approaches used for surfactin production.

<sup>a</sup> Maximum yield in whole culture. <sup>b</sup> Maximum productivity in whole culture. <sup>c</sup> Specific production when reaching maximum yield.

#### **4. Conclusions**

We investigated the effects of LDHs on the production of biomass and surfactin in a *B. subtilis* ATCC 21332 culture. The highest yield of surfactin (5280 mg/L) was obtained after 5 days of cultivation in the presence of 4 g/L Mg2Fe–LDH. This study demonstrated that LDHs have potential for use as additives to enhance the production of surfactin in *B. subtilis* ATCC 21332. Furthermore, microscopy revealed the inhibition of cell growth in the presence of the LDH, suggesting an efficient process for the production of surfactin through greater conversion of the carbon source.

**Supplementary Materials:** The following are available online at http://www.mdpi.com/2073-4352/9/7/355/s1, Figure S1: Typical surfactin standard chromatogram in HPLC showing surfactin isoform A–F.

**Author Contributions:** Data curation and methodology, P.-H.C. and S.-Y.L., conceptualization, writing—original draft preparation; writing—review and editing, project administration, T.-Y.J. and Y.-C.L.

**Funding:** This study was supported by research grants from the National Science Council of Taiwan, R.O.C. (grant nos. NSC 101-2221-E-005-061 and 106-2113-M-039-007), and China Medical University (grant no. CMU 107-N-22).

**Acknowledgments:** We thank Wei Yu-Hong of Yuan Ze University for sharing the strain *B. subtilis* ATCC 21332.

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


© 2019 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (http://creativecommons.org/licenses/by/4.0/).

*Article*

## **Rapid Removal and E**ffi**cient Recovery of Tetracycline Antibiotics in Aqueous Solution Using Layered Double Hydroxide Components in an In Situ-Adsorption Process**

#### **Kwanjira Panplado 1, Maliwan Subsadsana 2, Supalax Srijaranai <sup>1</sup> and Sira Sansuk 1,\***


Received: 9 June 2019; Accepted: 1 July 2019; Published: 4 July 2019

**Abstract:** This work demonstrates a simple approach for the efficient removal of tetracycline (TC) antibiotic from an aqueous solution. The in situ-adsorption removal method involved instant precipitation formation of mixed metal hydroxides (MMHs), which could immediately act as a sorbent for capturing TC from an aqueous solution, by employing layered double hydroxide (LDH) components including magnesium and aluminum ions in alkaline conditions. By using this approach, 100% removal of TC can be accomplished within 4 min under optimized conditions. The fast removal possibly resulted from an instantaneous adsorption of TC molecules onto the charged surface of MMHs via hydrogen bonding and electrostatically induced attraction. The results revealed that our removal technique was superior to the use of LDH as a sorbent in terms of both removal kinetics and efficiency. Moreover, the recovery of captured TC was tested under the influence of various common anions. It was found that 98% recovery could be simply achieved by using phosphate, possibly due to its highly charged density. Furthermore, this method was successful for efficient removal of TC in real environmental water samples.

**Keywords:** tetracycline; metal hydroxides; layered double hydroxides; removal; water sample

#### **1. Introduction**

The contamination of antibiotics in environmental water is a global concern, as they are potentially toxic and harmful to both the ecosystem and human life [1]. One of the most commonly found antibiotics is tetracycline (TC), widely used in a variety of animal livestocks in order to promote growth and kill bacteria [2]. It also plays a significant role in human therapy. As a result of its misuse, TC is commonly released as an agricultural or community effluent into the environment [3,4]. Furthermore, TC can be transferred from the community to a water source such as a lake. Therefore, the purification of wastewater by the removal of TC prior to release into environmental water is still necessary.

Even though treatment methods have been developed in wastewater management, most of them such as coagulation [5], ion exchange [6], and photocatalytic degradation [7] involve intensive energy and a high cost of operation, complex procedures, and production of possible toxic products. Above all reported methods, adsorption is the most reliable and effective means, with various sorbent materials being employed. For example, multi-walled carbon nanotubes (MWCNTs) as an adsorbent provided 99.8% removal of TC within 20 min [8]. High adsorption efficiency resulted from abundant π-π interactions found from the π systems on the MWCNT surface, and benzene rings and double bonds, both C=C and C=O, in TC molecules. Similarly, it was reported that graphene oxide functionalized

with magnetic particles gave 98% removal efficiency within 10 min as a consequence of strong π-π interactions between sorbent and target molecules [9]. Furthermore, zeolitic imidazolate metal organic framework ZIF-8 nanoparticles were used as an adsorbent and gave a removal efficiency of 90% as a result of a weak electrostatic interaction [10]. Furthermore, the removal of TC using Fe/Ni nanoparticle sorbent in an aqueous solution reached 97.4% removal within 2 h. [11]. In this sorbent-based adsorption technique, the removal efficiency of TC is chiefly dependent on the applied sorbents as well as their properties.

Among all these adsorbents, layered double hydroxide (LDH) has been considered as environmentally benign, and an effective adsorbent due to its high surface area and excellent anion exchangeability [12,13]. This material fundamentally consists of layers of mixed metal hydroxides linked with water molecules located in the interlayer space. The formula of LDH is generally symbolized as [*M*1−*<sup>x</sup>* <sup>2</sup>+*Mx* <sup>3</sup><sup>+</sup>(OH)2] *<sup>x</sup>*+(*An*−)*x*/*n*·mH2O, where *<sup>M</sup>*2<sup>+</sup> and *<sup>M</sup>*3<sup>+</sup> are divalent and trivalent cations, respectively. *x* is a molar ratio of *M*<sup>3</sup>+/(*M*2<sup>+</sup> + *M*3+) and *A* is an anion located at interlayers with the negative charge *n*. LDH has been employed as a sorbent or functionalized for the removal of various pollutants [14–21]. However, the synthesis and characterization of LDH sorbents are commonly required. Unavoidably, these steps are commonly time- and energy-consuming.

This study presents an in situ method for removal of TC antibiotics from an aqueous solution by employing the LDH components to generate instantly formed mixed metal hydroxides (MMHs), which simultaneously act as a sorbent to capture TC molecules from an aqueous solution during their precipitation formation. Unlike other sorbent-based removal techniques, our method was simple, efficient, rapid, and eco-friendly, as the synthesis of sorbent and its characterization can be avoided, thus saving time, energy, and cost. The parameters affecting the removal efficiency were investigated thoroughly. For comparison, a kinetic removal of this contaminant with LDH used as sorbent in a conventional route was also studied. The recovery of captured TC through an ionic interfering effect was demonstrated. Additionally, the removal of TC was carried out in real natural water samples.

#### **2. Materials and Methods**

#### *2.1. Chemicals and Reagents*

Tetracycline hydrochloride (C22H24N2O8·HCl) was obtained from Sigma-Aldrich (Hong Kong, China). Aluminium chloride hexahydrate (AlCl3·6H2O) and magnesium chloride hexahydrate (MgCl2·6H2O) were purchased from Sigma-Aldrich (St Louis, Missouri, USA). Sodium hydroxide (NaOH), sodium acetate (CH3COONa), sodium nitrate (NaNO3), and sodium sulphate (Na2SO4) were obtained from Carlo Erba (France). Sodium chloride (NaCl) and sodium carbonate (Na2CO3) were purchased from Ajax Finechem (Australia). Di-potassium hydrogen phosphate (HK2PO4) was obtained from BDH Prolabo (England). Methanol (CH3OH) was purchased from LiChrosolv®(Darmstadt, Germany). All chemicals and reagents were of at least analytical grade and used as received without further purification. Deionized (DI) water was used throughout. A 1000 mg L−<sup>1</sup> stock solution of TC was freshly prepared by dissolving 0.1000 g of C22H24N2O8·HCl in 100 mL of a mixed methanol:water solution (30:70% v/v). The solutions of mixed metal solutions (Mg2<sup>+</sup> and Al3<sup>+</sup>) with varied mole ratios were prepared by a dissolution of an appropriate amount of MgCl2·6H2O and AlCl3·6H2O in DI water.

#### *2.2. Synthesis of LDH*

As a comparative study, the removal of TC from an aqueous solution with LDH sorbent was investigated. First, LDH with a 3:1 mole ratio of Mg:Al was synthesized by co-precipitation as explained elsewhere [22]. Typically, a solution containing 0.0681 mol of MgCl2·6H2O (13.84 g) and 0.0227 mol of AlCl3·6H2O (5.48 g) in 200 mL DI was prepared. Then, 150 mL of an aqueous solution of 4.45 g NaCl was added slowly under stirring at room temperature and then 3 M NaOH was added to adjust the pH to 12. The mixture was then transferred into the Teflon coated stainless steel autoclave for hydrothermal treatment at 120 ◦C for 48 h. Next, the precipitates were filtered, washed thoroughly with DI water, and dried at 90 ◦C for 24 h.

#### *2.3. Measurement and Characterization*

To evaluate the removal efficiency, the absorption spectra of TC residuals were recorded by an Agilent 8453 UV-vis spectrophotometer (Agilent Technologies, Waldbronn, Germany). A centrifuge (H-11n, Kokusan, Tokyo) with 15 mL calibrated tubes was used for the phase separation. Power X-ray diffraction (XRD, D8 Advance, Bruker, Bremen, Germany) and attenuated total reflection Fourier transform infrared spectroscope (ATR-FTIR; Tensor 27, Bruker, Germany) were used. Scanning electron microscopy (SEM) was completed using a SNE-4500M microscope (SEC Co., Ltd, Seoul, South Korea). In addition, the surface charge of MMHs was assured by Zeta potential measurement using the Zetasizer Nano S (Malvern Instruments Ltd, Malvern, UK).

#### *2.4. Removal Procedure*

It can be noted that other divalent and trivalent cations such as Mn2<sup>+</sup>, Co2<sup>+</sup>, Ni2<sup>+</sup>, Zn2<sup>+</sup>, and Fe3<sup>+</sup> can be employed [23–25]. However, the rate of precipitation of metal hydroxides should be considered as it can result in the kinetic removal of TC. From the cost and toxicity point of view, Mg2<sup>+</sup> and Al3<sup>+</sup> ions were selected in this study. The removal of TC molecules from an aqueous solution was studied in a batch experiment. First, a simulated water sample containing 30 mg L−<sup>1</sup> TC was prepared and then conditioned with 1M NaOH. To initiate the removal of TC, a solution of Mg<sup>2</sup>+/Al3<sup>+</sup> ions with a fixed mole ratio of 3:1 was then added into 10 mL of TC solution. After a certain contact time, the TC-MMH particles were separated from the solution by centrifuging at 6000 rpm for 2 min. The residual TC in the solution was monitored by UV-vis measurement. The absorbance at 384 nm was used to evaluate the removal efficiency (%R) based on the following expression:

$$\% \mathbb{R} = [(A\_0 - A)/A\_0] \times 100$$

where *A*<sup>0</sup> and *A* are the absorbance of TC before and after the removal at any time, respectively. The effects of all factors on the removal efficiency including the volume of the mixed metal solution, the volume of NaOH, mole ratio of Mg2+/Al3+, and the removal time were studied.

For a comparative study, the kinetic removal of TC by the LDH sorbent was investigated. The experiment was performed in a batch system with a total volume of 10 mL of 30 mg L−<sup>1</sup> of TC and 0.0131 g of LDH sorbent. After a certain contact time from 15 s to 150 min, 1 mL of the solution was withdrawn for UV-vis measurement and 1 mL of DI water was introduced into the studied system. The removal efficiency was calculated by using the above expression.

#### **3. Results and Discussion**

#### *3.1. Removal of TC*

In this study, the in situ-adsorption removal strategy was based on the usage of LDH components in order to generate the mixed metal hydroxides (MMHs), which can simultaneously capture the TC molecules during their precipitation formation. First, the removal of TC from an aqueous solution was simply tested by introducing a metal solution containing both Mg2<sup>+</sup> and Al3<sup>+</sup> ions into an alkaline TC solution. In general, regarding its p*Ka* values (3.3, 7.7, and 9.7), TC is present in the aqueous solution in various forms depending on the pH of solution [26]. TC is in the TCH2 neutral form when the pH falls in the range of 3.3–7.7. However, it is protonated and in a positively-charged TCH3 <sup>+</sup> form in a solution with a pH less than 3.3 and it is negatively-charged as the TCH− form in alkaline conditions (pH > 10). As our strategy was performed in an alkaline solution at a pH of about 12, a removal of TC from an aqueous solution was accomplished through an interaction with instantly formed MMH acting as a sorbent. The removal resulted from an adsorption of TC molecules on the surface of MMHs as a consequence of an electrostatic interaction and H-bonding [27]. The UV-vis spectra of 30 mg L−<sup>1</sup> TC before and after the removal for 4 min of contact time is presented in Figure 1. It was found that the absorbance of TC was reduced tremendously after the removal. This indicated the potential of the present approach for TC removal.

**Figure 1.** UV-vis absorption spectra of TC in aqueous solution at an initial concentration of 30 mg L<sup>−</sup>1; before and after removal for 4 min contact time, with corresponding photograph images of the solution (inset).

#### *3.2. Optimization of the Removal of TC*

This method involved the use of LDH components to initiate the removal of TC from an aqueous solution. Accordingly, the main parameters including the amount of hydroxyl and metal ions were studied. These precursors could have an impact on the precipitation formation of MMHs, thereby affecting the removal efficiency. First, the impact of OH- on the removal of 30 mg L−<sup>1</sup> TC was investigated by varying the amount of OH- , while other factors were fixed. Figure 2a shows that the removal efficiency of TC was enhanced rapidly, when increasing the amount of OH- up to 200 mmol (200 mL). Then, the removal slightly improved until an addition of 275 mmol (275 mL) OH−. After that, a further addition of OH− did not benefit the removal of TC molecules. Thus, the amount of 275 mmol (275 mL) of OH- was optimal.

**Figure 2.** The effect of (**a**) the volume of 1 M NaOH and (**b**) the volume of mixed metal solution on the removal efficiency of TC in aqueous solution at an initial concentration of 30 mg L<sup>−</sup>1.

The effect of both Mg2<sup>+</sup> and Al3<sup>+</sup> ions on the removal of TC was also examined. The mole ratio of Mg2<sup>+</sup>:Al3<sup>+</sup> was fixed at 3:1 as an optimal ratio, while the volume of these ions was varied and other parameters were kept constant. As presented in Figure 2b, the removal efficiency of TC was improved with increasing the volume of the mixed metal solution up to 200 μL. After that, the removal of TC became slightly lower. Similar to OH−, an increase of Mg2<sup>+</sup> and Al3<sup>+</sup> content enhanced the precipitation formation of MMHs, thereby improving the interaction with TC in the solution. Thus, the removal efficiency increased when the components used to produce MMH sorbent increased as the removal is commonly dependent on the content of sorbent.

The mole ratio of Mg2+:Al3<sup>+</sup> was also investigated. In general, LDHs are synthesized by the hydrothermal method at a 2:1 and 3:1 mole ratio of Mg2+:Al3+. In this study, we simulated the formation of LDHs and employed their components to generate the simultaneous removal of TC from the solution. Under our studied conditions, the mole ratios of these ions varied from 1:1 to 4:1, while other parameters were kept constant. As can be seen in Figure 3a, the TC removal efficiency was sharply boosted by extending the mole ratio of Mg2+:Al3<sup>+</sup> ions up to 2:1. It was noticed that the strong precipitation of MMHs probably depended on the content of Mg2<sup>+</sup> more than Al3+. This resulted in an increase in the removal of TC when the content of Mg2<sup>+</sup> ions was increased. However, the removal of TC reached the maximum value with the Mg2<sup>+</sup>:Al3<sup>+</sup> mole ratio of 3:1. Therefore, Mg2+:Al3<sup>+</sup> with a mole ratio of 3:1 was appropriate for the removal of TC by this approach.

The kinetic removal of TC by the proposed method was studied. The removal experiments were carried out at various contact times. The contact time was described as the time required for TC removal after the introduction of a mixed metal solution into the studied system. The experiments were tested under an optimization condition including 275 μL of NaOH, 200 μL of Mg2<sup>+</sup>:Al3<sup>+</sup> at a 3:1 mole ratio and an initial TC concentration of 30 mg L−1. Figure 3b shows the kinetic removal of TC by our method. It was observed that the removal efficiency of TC increased rapidly and slightly improved after 15 s. However, almost complete removal of TC was achieved after 4 min of contact time. The comparison of the removal of TC by the proposed method with other methods is summarized in Table 1. The results indicated that the proposed technique was more efficient and rapid than other removal methods. In addition, our method required no synthetic step of sorbents or catalysts, thus saving time, chemicals and energy.

**Figure 3.** The effect of (**a**) mixed metal concentrations and (**b**) contact time on the removal efficiency of TC in aqueous solution at an initial concentration of 30 mg L<sup>−</sup>1.


**Table 1.** Comparison of the removal of TC by the proposed method with other methods.

#### *3.3. Characterization of MMH Sorbent and Confirmation for TC Removal*

To confirm the removal of TC by our strategy, an instantly formed MMH sorbent was collected and then characterized. First, the surface charge of MMH was confirmed by the Zeta potential measurement. The result indicated that the surface of MMH particles was positively-charged with +19 mV. In the presence of TC, it was observed that the surface charge of MMH particles was reduced to −0.1 mV. The reduction of surface charge possibly resulted from the adsorption of TC molecules, present as a negative form (TCH−) in the studied alkaline conditions, on the surface of MMH. The interaction of MMH with TC was also investigated by FT-IR measurement. As presented in Figure 4, when compared with the spectra of TC and MMH sorbent (without TC removal), a decrease in the intensity of the sharp peak at approximately 3700 cm<sup>−</sup>1, corresponding to the -OH group in the MMH sorbent after removal of TC, was observed. This possibly implied a surface interaction of MMH with TC molecules via hydrogen bonding. It was also found that after the removal of TC, a characteristic peak at 1524 cm−1, which was attributed to the vibration of the NH2 group of TC molecules, shifted to 1597 cm<sup>−</sup>1. These results confirmed the interaction between TC and MMH particles instantly generated in the system. Moreover, the presence of TC on the surface of MMH was also investigated by SEM measurement. As displayed in Figure 5, it was revealed that TC appeared on the surface of the MMH sorbent, confirming the interaction of TC molecules with MMH sorbent.

**Figure 4.** FT-IR spectra of pure TC, MMHs, and TC-MMHs.

**Figure 5.** SEM images of (**a**) MMHs and (**b**) TC-MMHs.

#### *3.4. Comparison on TC Removal with LDH Sorbent*

We also investigated the removal of TC by using an LDH sorbent. LDH was prepared by a hydrothermal method, as explained in Section 2.2. Figure 6a represents the SEM image of LDH. It was seen that typical plate-like particles were observed. In addition, the XRD pattern confirmed the characteristics of LDH (see Figure S1 in Supplementary Materials). The experiments were conducted in a traditional way as a batch system at 30 mg L−<sup>1</sup> of TC solution, as described in Section 2.4. An amount of LDH sorbent was calculated based on the optimized removal condition of the proposed method. LDH (0.0131 g) was obtained and further employed as a solid sorbent. The removal kinetics of LDH sorbent, in comparison with the proposed strategy, are displayed in Figure 6b. It was clearly seen that the removal of TC by LDH sorbent was quite slow. The removal reached an equilibrium after 60 min with a removal efficiency of 75%. This result indicated an incomplete removal of 30 mg L−<sup>1</sup> of TC by the LDH sorbent. It is noted that the removal efficiency can be improved with increasing the amount of sorbent. However, this removal kinetic result implied that the present method was superior to the traditional sorbent-based removal method in terms of rapidness, efficiency, and cost-effectiveness as it eliminated the synthesis step of sorbent, thereby saving time, chemicals, and energy.

**Figure 6.** (**a**) SEM image of LDH sorbent and (**b**) its kinetic removal of TC in aqueous solution at an initial concentration of 30 mg L<sup>−</sup>1, compared with MMHs.

#### *3.5. Recovery of Captured TC*

The ability of the proposed removal approach to recover TC from an aqueous solution was also tested. After the removal process, the precipitates containing TC were dispersed for 1 min in a solution containing an excess of various common anions (sodium salts) including Cl−, NO3 <sup>−</sup>, CH3COO−, SO4 <sup>2</sup>−, CO3 <sup>2</sup>−, and PO4 <sup>3</sup>−. After centrifugation for phase separation, the solution was then taken for analysis of TC content. These anions were employed as they could interfere with the adsorption ability of TC molecules on the surface of MMH particles. Thus, recovery of TC could be achieved. Figure 7 represents the UV-vis spectra of TC collected after the recovery step and the recovery efficiency regarding different anions used. It was found that the recoveries of TC were ordered with the use of PO4 <sup>3</sup><sup>−</sup>, CO3 <sup>2</sup><sup>−</sup>, SO4 <sup>2</sup><sup>−</sup>, NO3 <sup>−</sup>, CH3COO<sup>−</sup>, and Cl−. The effect of these anions on both the precipitation formation of MMH particles and adsorption of TC possibly resulted from their charge density. Hence, 98% of recovery was obtained in the case of PO4 <sup>3</sup><sup>−</sup> used. In addition, this exhibited the advantage of the present removal strategy as the recovery of TC was simply obtained, regarding the loosely formed MMH sorbent.

**Figure 7.** (**a**) UV-vis absorption spectra of TC and (**b**) the recovery efficiency after recovery of TC with various anions.

#### *3.6. Removal of TC in Real Environmental Water Samples*

The applicability of the proposed method was further assessed by the removal of TC in environmental water samples. The water samples were collected from various water sources including Nam Pong River, Ubonlratana Dam, Nong Kot Lake, Si Than Lake, Kaennakorn Lake, and the wastewater treatment pond located at Srinagarind Hospital, Khon Kaen, Thailand. It was found from an initial analysis that these water samples had a pH in the range of 6.6–7.5 without the presence of TC. Accordingly, to evaluate the applicability of the proposed method, all water samples were first filtrated and then spiked with 30 mg L−<sup>1</sup> of TC. After removal using our strategy, the removal efficiency was reported, as presented in Figure 8. The results revealed that the removal of TC performed in every simulated water sample was 92%. The result indicated the applicability of our approach for the removal of TC in real water samples. High removal efficiencies were obtained even with the presence of other ionic species or contaminants, commonly found in natural water. Therefore, the interference of other compounds or ions is negligible. It can also be noted that other ionic species could facilitate the precipitation formation of MMHs, increasing the adsorption ability of TC. This led to an enhanced removal efficiency of TC from the aqueous solution.

**Figure 8.** Removal of TC added in various environmental water samples at an initial concentration of 30 mg L<sup>−</sup>1.

#### **4. Conclusions**

This study demonstrates a simple strategy for the efficient removal of tetracycline (TC) antibiotic from an aqueous solution. The in situ-adsorption method involves the utilization of layered double hydroxide (LDH) components to initiate precipitation of mixed metal hydroxides (MMHs), concurrently acting as a sorbent for instant adsorption of TC molecules from an alkaline solution. Mg2<sup>+</sup> and Al3+, present as chloride forms, were used. Under optimized conditions, 99.5% removal efficiency can be obtained within 4 min due to a strong electrostatic interaction and hydrogen bonding between TC molecules and the positively-charged surface of MMHs. When compared with the removal adsorption by the LDH sorbent, our method is much better and faster. With our removal strategy, 98% recovery of TC captured by MMHs can be simply achieved by dispersion in a phosphate solution. Moreover, almost complete removal of TC from simulated environmental water samples including a dam, river, three lakes and wastewater plant can be obtained.

**Supplementary Materials:** The following are available online at http://www.mdpi.com/2073-4352/9/7/342/s1, Figure S1: XRD pattern of LDH

**Author Contributions:** Conceptualization, S.S. (Sira Sansuk); investigation, K.P. and M.S.; methodology, K.P. and M.S.; supervision, S.S. (Supalax Srijaranai); writing—original draft preparation, K.P.; writing—review and editing, S.S. (Sira Sansuk).

**Funding:** This research received no external funding.

**Acknowledgments:** We gratefully acknowledge the Science Achievement Scholarship of Thailand (SAST).

**Conflicts of Interest:** The authors declare no conflict of interest.

#### **References**


© 2019 by the authors. Licensee MDPI, Basel, Switzerland. This article is an open access article distributed under the terms and conditions of the Creative Commons Attribution (CC BY) license (http://creativecommons.org/licenses/by/4.0/).
