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Abstract

In the present work, natural diatomite modified with manganese oxide (MnO₂) was prepared via direct redox reaction with KMnO₄ and HCl. The product was characterized by using X-ray diffraction, Fourier-transform infrared spectroscopy, scanning electron microscopy, transmission electron microscopy, energy-dispersive X-ray, and nitrogen adsorption-desorption isotherms. It was found that the nanorod manganese oxide was highly dispersed onto the diatomite porous matrix. The specific surface area of the obtained manganese oxide/diatomite ( $S_{BET} = 68.5 m^{2} g^{- 1}$ ) is larger than that of natural diatomite ( $S_{BET} = 55.4 m^{2} g^{- 1}$ ). It was utilized to remove Pb(II) in aqueous solutions. It exhibits an excellent Pb(II) adsorption capacity. The adsorption data fits well with the pseudo-second-order kinetics model, and the adsorption process is endothermic and spontaneous with an activation energy of 41.56 kJ mol⁻¹ and follows the Freundlich isotherm model. The Mn/diatomite adsorption capacity for Pb(II) is 81.42 mg g⁻¹, calculated with the Langmuir model. In addition, the adsorption mechanism of Pb(II) onto Mn/diatomite is also addressed.

1. Introduction

Heavy metals are commonly found in wastewater from the chemical industry, paints, ceramics, glass, mining, and battery manufacturing and severely harm humans and ecosystems [1, 2]. Among them, lead (Pb(II)) has an amphoteric structure and is a metal with a high toxic level among heavy metals. Pb(II) is easily dispersed through the soil and water ecosystem, and by entering the food chain, it causes toxic effects on human health [3]. Due to exposure to Pb(II) for a long period, the kidneys and immune systems, in particular, can be severely destroyed. Therefore, various methods have been utilized to remove heavy metal ions from aqueous solutions, such as oxidation, coagulation and precipitation, membrane filtration, adsorption, ion exchange, and biological treatment [4–9]. Among them, adsorption is the most effective because of its simplicity, low cost, and ability to treat wastewater containing low concentrations of metal ions [10–14].

Diatomite (SiO₂·nH₂O) is a soft, grey, low-density sedimentary rock. The main component of diatomite is the shell of diatoms consisting primarily of amorphous silica. Diatomite has unique properties, such as high porosity, high adsorption capacity, and high heat resistance [15–19]. These properties suggest that diatomite is a potential adsorbent for the pollutants present in industrial wastewater. Furthermore, natural diatomite is abundant, cheap, and environmentally friendly [18]. In Vietnam, diatomite mineral is often applied in shrimp pond treatment, sound and heat insulation, etc. Recently, the modification of diatomite to increase its applicability has also attracted scientists [20–24]. The modification via hydrothermal process exhibits advantages including nanoparticles with the desired size and shape [25] and formation of a well-crystallized powder [26]. Hence, these techniques have been widely applied to introduce the inorganic compounds with active sites to expand the application of diatomite in adsorption and catalysis [12, 27–30].

Nanoscale manganese oxide (MnO₂) particles, known to be a porous material and a mild oxidant, possess a large surface area and excellent chemical stability in acidic media [31]. Manganese oxide has been widely employed as a constituent material in air batteries [32] and a catalyst for oxygen reduction [33]. In particular, manganese oxide exhibits excellent adsorption toward heavy metals in aqueous sources. However, nanosized MnO₂ particles suffer from heavy agglomeration, limiting their application [34]. Moreover, using pure MnO₂ as an adsorbent is not reasonable regarding economic effectiveness, and it is very difficult to separate the particles from the liquid phase [35]. To deal with these limitations and to promote the treatment efficiency and capabilities for eliminating heavy metals in an aqueous solution with natural diatomite, numerous researchers modified diatomite by coating nanosized MnO₂ particles onto its surface. Al-Degs et al. [10] and Khraished et al. [11] modified diatomite with manganese oxide by treating it with manganese chloride and sodium hydroxide. The results demonstrate that its heavy metal adsorption capacity was much improved compared with that of the original diatomite. The specific surface area of diatomite and Mn/diatomite calculated with the simple and rapid methods was 33 and 80 m² g⁻¹. The adsorption capacity of diatomite is 24.9, 27.6, and 16.1 mg g⁻¹ adsorbates for Pb²⁺, Cu²⁺, and Cd²⁺, respectively, while the corresponding values of Mn/diatomite are 99.0, 57.6, and 27.9 mg g⁻¹. Li et al. [36] reported a two-step procedure for diatomite modification with nano-MnO₂, and the material has a maximum adsorption capacity of 56.84 mg g⁻¹ for Pb²⁺. Du et al. [12] also successfully synthesized MnO₂ nanowires on a diatomite substrate. In addition, MnO₂-modified diatomite was also used for Zn(II) adsorption [13], dye adsorption [37], catalysts [38], and electrode modifiers [39].

In this paper, a manganese-modified diatomite material (Mn/diatomite) was prepared with the hydrothermal method via the redox reaction of KMnO₄ and HCl, where diatomite was a substrate. The material’s physicochemical properties and ability to remove Pb(II) ions from aqueous solutions were investigated.

2. Experimental

2.1. Materials and Chemicals

Raw diatomite was taken from Phu Yen province, Vietnam, washed several times with water, filtered, dried at 100°C, sieved, and stored in closed containers for further tests. KMnO₄ (Merck, Germany), Pb(CH₃COO)₂·3H₂O, NaOH, HCl, and KCl (Guangdong, China) are of analytical grade and are used without further purification.

2.2. Preparation of Mn/Diatomite

Diatomite was modified with manganese, according to Wu et al. [40] and Du et al. [12]. Briefly, (i) 0.5 g of diatomite, 5 mmol of KMnO₄ (0.79 g), and 80 mL of distilled water were mixed to obtain mixture A; (ii) a certain volume of a 10 M HCl solution was added to mixture A under stirring to obtain mixture B; (iii) the entire mixture B was transferred into a 200 mL Teflon-lined steel autoclave with a quantity of distilled water to fill 80% of the Teflon-flask volume to obtain mixture C; (iv) the Teflon flask was placed in an oven at a specified temperature and for a certain time; and (v) the flask was cooled to ambient temperature, and the solid product was collected, filtered, washed with distilled water, and dried at 60°C for 24 h. The resulting solid is a modified diatomite material and is denoted as Mn/diatomite.

The effects of reaction conditions on the modification, including hydrothermal temperature (80, 120, and 160°C), hydrothermal time (16, 24, and 48 h), and the KMnO₄/HCl molar ratio (1 : 1, 1 : 2, 1 : 4, and 1 : 8) were investigated. The details of the experimental conditions used in the diatomite modification are summarized in Table 1.

Table 1

Experimental conditions were carried out during the preparation of the Mn/diatomite samples.

Sample	Mass of diatomite (g)	Mass of KMnO₄ (g)	10 M HCl volume (mL)	KMnO₄/HCl molar ratio	Hydrothermal temperature (°C)	Hydrothermal time (h)
1	0.5	0.79	4	1 : 8	80	16
2	0.5	0.79	4	1 : 8	120	16
3	0.5	0.79	4	1 : 8	160	16
4	0.5	0.79	0.5	1 : 1	80	16
5	0.5	0.79	1	1 : 2	80	16
6	0.5	0.79	2	1 : 4	80	16
7	0.5	0.79	2	1 : 4	80	24
8	0.5	0.79	2	1 : 4	80	48

2.3. Preparation of Nano-MnO₂

Nano-manganese oxide was prepared according to a similar procedure in Section 2.2. The 10 M HCl solution was added drop-wise (the molar ratio of KMnO₄/HCl was adjusted in the ratio of 1 : 1, 1 : 2, 1 : 4, and 1 : 8). The Teflon flask was heated at 160°C for 16 h, and the yielded nano-manganese oxide was dried at 60°C.

2.4. Characterization

The morphology of the materials was observed by using scanning electron microscopy (SEM, IMS-NKL) with an acceleration voltage of 5 kV and transmission electron microscopy (TEM, EMLab-NIHE) at a voltage of 80 kV. X-ray diffraction (XRD) patterns were recorded on a VNU-D8 Advance Bruker (Germany) powder diffractometer with a Cu Kα radiation source and an angle and time scan step of 0.03° and 0.04 s, respectively. Fourier-transform infrared (FT-IR) spectra were obtained on a Jasco FT/IR-4600 (Japan), with samples being dispersed on KBr pallets and measured within the wavenumber range of 4000-400 cm^-1 with a spectral resolution of 1 cm^-1. Nitrogen adsorption measurements were conducted with a Micromeritics Tristar 3000 apparatus at the liquid nitrogen temperature (77 K). The textural properties of the material were examined via the specific surface area and porosity after a heat treatment at 250°C in the N₂ for 5 h. The elemental composition on the material surface was analyzed with energy-dispersive X-ray (EDX) spectroscopy along with the SEM (JEOL JED-2300) at 20 kV. The Pb(II) concentration was determined with the atomic absorption spectroscopy (AAS) method on a Shimadzu AA-7000 (Singapore).

2.5. Adsorption Experiments

2.5.1. Evaluate the Pb(II) Adsorption Capacity of the Material

A quantity of adsorbent (0.1 g) was placed into an Erlenmeyer flask containing 100 mL of a Pb(II) solution with a prescribed concentration and shaken at ambient temperature for 3 h to achieve adsorption-desorption equilibrium. Then, the solution was filtered to remove the adsorbent, and the Pb(II) ion concentration in the solution was determined.

The adsorption capacity at equilibrium ( $q_{e}$ ) was calculated according to the following equation. $\begin{matrix} (1) & q_{e} = \frac{(C_{0} - C_{e}) \times V}{m}, \end{matrix}$ where $C_{0}$ and $C_{e}$ (mg L^-1) are the concentration of Pb(II) in the solution at the beginning and at equilibrium, respectively; $V$ (L) is the volume of the Pb(II) solution; and $m$ (g) is the weight of Mn/diatomite used for each adsorption.

2.5.2. Pb(II) Adsorption Kinetics

Mn/diatomite (0.1 g) was placed into a double-necked flask containing 100 mL of the Pb(II) solution with a prescribed concentration under magnetic stirring at a prescribed temperature, under natural pH. At a specified interval, a certain volume of solution was withdrawn and centrifuged to remove the adsorbent, and the Pb(II) concentration in the supernatant was determined with the AAS technique.

The adsorption capacity at time $t$ ( $q_{t}$ ) was calculated according to the following equation. $\begin{matrix} (2) & q_{t} = \frac{(C_{0} - C_{t}) \times V}{m}, \end{matrix}$ where $C_{t}$ (mg L^-1) is the concentration of Pb(II) in the solution at time $t$ .

The pseudo-first-order (Eq. (3)) and pseudo-second-order (Eq. (4)) kinetic models are commonly used to test experimental data [13, 41–47]. $\begin{matrix} (3) & q_{t} = q_{1} \times (1 - e^{- k_{1} \times t}), \\ (4) & q_{t} = \frac{q_{2}^{2} \times k_{2} \times t}{(1 + q_{2} \times k_{2} \times t)}, \end{matrix}$ where $k_{1}$ (min^-1) and $k_{2}$ (g mg^-1 min^-1) and $q_{1}$ and $q_{2}$ (mg g^-1) are the rate constants and the maximum adsorption capacities of the pseudo-first-order and pseudo-second-order adsorption kinetic models, respectively.

2.5.3. Effect of Solution pH

Mn/diatomite (0.02 g) was placed into an Erlenmeyer flask containing 50 mL of the Pb(II) solution with a specified concentration and pH (the pH value of the solution was adjusted with a 0.1 M HCl or 0.1 M NaOH solution) and shaken at ambient temperature for 3 h to achieve adsorption-desorption equilibrium. Then, the adsorbent was removed, and the Pb(II) concentration was determined.

2.5.4. Equilibrium Studies

In this section, the adsorption was performed at 303 K, and the initial concentration of Pb(II) varied between 30.8 and 188.3 mg L^-1 ( $V = 0.1 L$ ; $m = 0.1 g$ ). To analyze the data of Pb(II) adsorption onto Mn/diatomite, we used the Freundlich, Langmuir and Sips models.

The Freundlich model assumes heterogeneous energy distribution of the active sites on the adsorbent surface with the interactions within the adsorbate. The Freundlich isotherm can be expressed in the following equation [13, 37, 41, 42, 45–47]. $\begin{matrix} (5) & q_{e} = K_{F} \times C_{e}^{1 / n}, \end{matrix}$ where $K_{F}$ is the Freundlich constant (mg^(1-1/n) L^1/n g^-1) and $n$ is the adsorption intensity. Theoretically, the adsorption conditions are favourable if the $n$ constant is greater than 1.

The Langmuir isotherm is a theoretical model for monolayer adsorption to a surface containing a finite number of adsorption sites with uniform adsorption energies without the displacement of the adsorbate in the plane of the adsorbent surface. The Langmuir isotherm is expressed in the following equation [10, 11, 37, 41–43, 45–47]. $\begin{matrix} (6) & q_{e} = \frac{q_{m} \times K_{L} \times C_{e}}{1 + K_{L} \times C_{e}}, \end{matrix}$ where the $K_{L}$ constant is related to the adsorption energy (L g^-1) and $q_{m}$ is the Langmuir monolayer adsorption capacity (mg g^-1).

The Sips isotherm has a form similar to the Freundlich isotherm, but it has a finite limit when the concentration is sufficiently high. The Sips equation can be expressed in the following equation [48]. $\begin{matrix} (7) & q_{e} = \frac{q_{m_{S}} \times K_{S} \times C_{e}^{m_{S}}}{1 + K_{S} \times C_{e}^{m_{S}}}, \end{matrix}$ where the $K_{S}$ (L ^m mg^m-1) and $q_{m_{S}}$ (mg g^-1) are the Sips equilibrium constant and maximum adsorption capacity, respectively, and $m_{S}$ is the Sips model exponent.

Isothermal and kinetic parameters were determined from nonlinear regression with the Solver tool in Microsoft Excel [49–51] with the root-mean-square error (RMSE) function described in the following equation [49, 50]. $\begin{matrix} (8) & RMSE = \sqrt{\frac{1}{N} \sum_{i = 1}^{N} {(q_{e, \exp} - q_{e, cal})}_{i}^{2}}, \end{matrix}$ where $q_{e, \exp}$ and $q_{e, cal}$ are the adsorption capacity values obtained from the experiment and models, respectively, and $N$ represents the number of data points.

2.5.5. Thermodynamic Studies

In this section, the initial concentration of Pb(II) was kept constant ( $V = 0.1 L$ , $m = 0.1 g$ , $C_{0} = 75 mg L^{- 1}$ ). The adsorption was performed at temperatures ranging from 303 to 338 K. The feasibility of the adsorption process can be estimated through thermodynamic studies. The thermodynamic parameters of Pb(II) adsorption on Mn/diatomite were calculated using the following equations [15, 37]. $\begin{matrix} (9) & K_{C} = \frac{C_{0} - C_{e}}{C_{e}} \times \frac{V}{m} = \frac{q_{e}}{C_{e}}, \\ (10) & \ln K_{C} = - \frac{Δ H °}{R \times T} + \frac{Δ S °}{R}, \end{matrix}$ where $R$ (8.314 J mol^-1 K^-1) is the ideal gas constant and $T$ (K) is the absolute temperature. $Δ H$ ° is the enthalpy change, and $Δ S$ ° is the entropy change of the adsorption process.

The free energy change $Δ G$ ° of the adsorption process is calculated according to the following equation [15, 37]. $\begin{matrix} (11) & Δ G ° = Δ H ° - T \times Δ S ° . \end{matrix}$

The adsorption rate constant ( $k$ ) of the adsorption process is a temperature-dependent function expressed in the following Arrhenius equation [41, 46]. $\begin{matrix} (12) & \ln k = \ln A - \frac{E_{a}}{R \times T}, \end{matrix}$ where $E_{a}$ is the Arrhenius activation energy (J mol^-1) and $A$ is the Arrhenius constant.

2.5.6. Desorption Experiments

After each experiment, the adsorbent was recycled. The desorption was performed by mixing the used adsorbent with 0.1 M HCl solution (a ratio of 0.5 g/50 mL). The mixture was then shaken with a shaker at ambient temperature ( $29 \pm 2$ °C) at 200 rpm for 120 min, filtered to recover the adsorbent, and washed with distilled water for subsequent readsorption.

3. Results and Discussion

3.1. Modification of Diatomite with Manganese Oxide and Pb(II) Adsorption Capacity Assessment

3.1.1. Effect of Hydrothermal Temperature

The morphology of natural diatomite and the Mn/diatomite samples prepared at different hydrothermal temperatures is presented in Figure 1. Natural diatomite exists as cylindrical tubes with a 5-7 μm diameter. Its surface is relatively smooth and has pores with a diameter of about 0.5 μm (Figure 1(a)). The morphology of manganese oxide varies from a flower-like structure to a needle-like one with an increase in the hydrothermal temperature. The Mn/diatomite sample prepared at 80°C exhibits flower-like particles attached to the diatomite surface and pores (Figure 1(b)). The morphology of the flower-like particle with some hundred nanometers in diameter consists of thin plates embroiled together. At the hydrothermal temperatures of 120 and 160°C, needle-like manganese oxide appears and is dispersed on the surface and in the pores of natural diatomite (Figures 1(c) and 1(d)).

Figure 1

SEM images of diatomite (a) and the Mn/diatomite samples prepared at different hydrothermal temperatures: (b) 80°C; (c) 120°C; (d) 160°C, and the insets present TEM images (experimental conditions: hydrothermal time of 16 h and molar ratio of KMnO₄/HCl 1 : 8).

(b)

(c)

(d)

Natural diatomite has a moderate adsorption capacity (15.6 mg g^-1) compared with Mn/diatomite (26.3-32.9 mg g^-1). Although all samples have a higher adsorption capacity, the one prepared at 80°C outweighs the others (32.9 mg g^-1) despite its lowest manganese content (Table S1). Thus, it can be assumed that manganese does not support Pb(II) adsorption. Instead, the material structure may favour adsorption because the denaturation at 80°C forms a material with a flower-like configuration. Therefore, this temperature was chosen for further experiments.

3.1.2. Effect of Molar Ratio of Precursors

The molar ratio between KMnO₄ and HCl was chosen at 1 : 1, 1 : 2, and 1 : 4. At the ratio of 1 : 1 (Figures 2(a) and 2(b)), the sample has a large number of circular or rod-shaped clusters attached to the surface of diatomite. At the ratio of 1 : 2 (Figures 2(c) and 2(d)), the circular clusters become larger, and spheres with a flower-like membrane structure appear (like the Mn/diatomite sample prepared at the molar ratio of KMnO₄/HCl 1 : 8; Figure 1(b)). Finally, at the ratio of 1 : 4 (Figures 2(e) and 2(f)), besides the spheres with the flower-like membrane structure, film-like structures are formed uniformly on the surface of the diatomite, thus increasing the Pb(II) adsorption efficiency. In this case, the structure of Mn-modified diatomite decides the Pb(II) adsorption efficiency, and the ratio of 1 : 4 provides the highest efficiency (43.1 mg g^-1, Table S1). Therefore, this ratio was chosen for subsequent experiments.

Figure 2

SEM images of Mn/diatomite samples prepared at different KMnO₄/HCl molar ratios: (a, b) 1 : 1; (c, d) 1 : 2; (e, f) 1 : 4 (experimental conditions: hydrothermal time and temperature are 16 h and 80°C, respectively).

(b)

(c)

(d)

(e)

(f)

3.1.3. Effect of Hydrothermal Time

In this section, we investigate the hydrothermal time under the conditions found above (Figure 3). The figure shows the SEM images of the 24 and 48 h samples with flower-like spheres and a regular layer of membranous structures. This topography is similar to that of the 16 h samples (Figures 2(e) and 2(f)). Concerning the content of manganese, it seems that a longer hydrothermal time favours the amount of this element in the sample and thus, together with the membranous structures, enhances Pb(II) adsorption (Table S1). From these results, we choose 24 h (54 mg g^-1 Pb(II) adsorption efficiency) as the appropriate time for further studies. The 24 h sample was denoted as Mn/diatomite and used for characterization.

Figure 3

SEM images of Mn/diatomite samples prepared at different hydrothermal times: (a, b) 24 h; (c, d) 48 h (experimental conditions: hydrothermal temperature of 80°C and molar ratio of KMnO₄/HCl 1 : 4).

(b)

(c)

(d)

The crystalline phase of natural diatomite and the Mn/diatomite was analyzed based on the X-ray data. For natural diatomite, the XRD pattern displays characteristic diffraction peaks of amorphous silica (broad diffraction centered at 20-25°) [52–56]. As for Mn/diatomite samples prepared at different hydrothermal temperatures, no characteristic diffraction peaks of manganese oxide are observed (Figure 4(a)), which is probably because manganese oxide particles are very small and evenly distributed on the surface of diatomite or because diatomite hinders the X-ray diffraction of manganese oxide. This result is consistent with that reported by Li et al. [36].

Figure 4

(a) XRD patterns of natural diatomite and the Mn/diatomite samples prepared at different hydrothermal temperatures (experimental conditions: hydrothermal time of 16 h and molar ratio of KMnO₄/HCl 1 : 8); (b) XRD pattern of the Mn/diatomite sample (experimental conditions: hydrothermal time and temperature of 24 h and 80°C, respectively, and molar ratio of KMnO₄/HCl 1 : 4); (c) nitrogen adsorption-desorption isotherms; (d) pore-size distributions of diatomite and Mn/diatomite (experimental conditions: hydrothermal time and temperature of 24 h and 80°C, respectively, and molar ratio of KMnO₄/HCl 1 : 4).

(b)

(c)

(d)

The XRD pattern of Mn/diatomite (Figure 4(b)) shows characteristic diffraction peaks of amorphous silica at 20–25° [52–56] and the peaks for MnO₂ (JCPDS No. 00-024-0735) with low intensity because of the oxide’s small size and its uniform distribution on the surface of diatomite, as mentioned above.

The textural properties of natural diatomite and Mn/diatomite indicate the type II isotherm and a H3-type hysteresis loop (Figure 4(c)), corresponding to the existence of macropores of nonuniform size and/or shape [57]. Besides, the Mn/diatomite sample with significant condensation at high relative pressures ( $P / P ° ~ 1$ ) indicates the existence of a slit-pore structure because the pores formed between the membrane of the modified material.

The pore-size distribution curve (Figure 4(d)) indicates that natural diatomite has pores of about 4.3 nm in diameter, while the Mn/diatomite sample has smaller pores (3.9 nm). This decrease is caused by the manganese oxide film coated on the surface. However, the specific surface area of Mn/diatomite ( $S_{BET} = 68.5 m^{2} g^{- 1}$ ) is larger than that of natural diatomite ( $S_{BET} = 55.4 m^{2} g^{- 1}$ ) (Table S2) because of the space formed between the sheets of manganese oxide attached to the diatomite surface. However, this specific area difference is insignificant.

3.2. Characteristic Properties and Pb(II) Adsorption Capacity of Nano-MnO₂

Nano-MnO₂ was prepared from KMnO₄ and HCl with four molar ratios of 1 : 1, 1 : 2, 1 : 4, and 1 : 8. All the samples have a bar shape with sides of about 60 to 100 nm (Figure 5), very similar to what was prepared by Wu et al. [40]. Some bars have a hollow tip like a tube. Except for the 1 : 1 KMnO₄/HCl sample, which has the bars stacked together, the other three samples have more or less separate bars with an insignificant number of stacks.

Figure 5

SEM images of MnO₂ samples synthesized at different molar ratios of KMnO₄/HCl: (a) 1 : 1; (b) 1 : 2; (c) 1 : 4; (d) 1 : 8 (experimental conditions: hydrothermal time and temperature are 16 h and 160°C, respectively).

(b)

(c)

(d)

The MnO₂ samples have the crystalline structure of an α-MnO₂ crystal with characteristic XRD peaks at 28.8, 37.6, 42.1, 49.8, 56.0, and 60.2° (Figure 6), corresponding to the (310), (211), (301), (411), (600), and (521) planes (JCPDS No. 00-024-0735) [40, 58, 59]. The sharp peaks with high intensity of the 1 : 2 to 1 : 8 KMnO₄/HCl samples indicate that the formed α-MnO₂ has high crystallinity and order. The sample at the KMnO₄/HCl 1 : 1 ratio has sharp peaks with low intensity, showing a low-ordered material. This observation is completely consistent with the results observed with SEM in Figure 5(a).

Figure 6

XRD pattern of MnO₂ samples synthesized at different KMnO₄/HCl molar ratios (experimental conditions: hydrothermal time and temperature are 16 h and 160°C, respectively).

The nano-MnO₂ samples were subjected to Pb(II) adsorption in the aqueous solution. Their adsorption capacities are similar with low values (16.2-18.9 mg g^-1 at $C_{0, Pb (II)} = 70.6 mg L^{- 1}$ ) (Figure S1). However, the Pb atoms are uniformly distributed over the entire area of the MnO₂ nanomaterial (Figure S2). The distribution of elements on the Mn/diatomite samples before Pb(II) adsorption shows that the material contains primarily Si and O with some Fe, Al, and Mn (Figure S3). After adsorption, the mapping displays Pb besides other elements. Manganese and Pb are uniformly distributed throughout the material surface, indicating that Pb adsorbs onto the Mn/diatomite material (Figure S4).

The FT-IR spectra of natural diatomite, Mn/diatomite, and Mn/diatomite after Pb(II) adsorption are relatively similar (Figure 7(a)), with a broad band at 1100 cm^-1 and two narrow bands at 797 and 470 cm^-1, corresponding to the asymmetric and symmetric stretching vibrations and bending vibration of the Si-O-Si bonds, respectively [13, 19, 55]. The peaks at 3697 and 3621 cm^-1 are typical for surface hydroxyl groups in diatomite, with the former corresponding to the isolated hydroxyl group (Si-OH) on the surface of diatomite [55, 60, 61], and the latter due to O-H stretching vibration of the aluminol groups (≡AlOH) [61]. These bands of Mn/diatomite have a lower intensity, indicating the chemical interaction between the surface silanol groups and the oxides [13]. The broad bands at 3441 and 1637 cm^-1 are thought to be due to the vibrations of the O-H bonds of adsorbed H₂O molecules, including free adsorbed water and hydrogen-bonded water with surface hydroxyl groups. The absorption band at 533 cm^-1 is thought to be the stretching vibration of the Fe-O bond [55] or the Me-O bond (with Me being a metal). The bands observed at 431-435 and 419-426 cm^-1 in Mn/diatomite and nano-MnO₂ (Figure 7(b)) are probably another vibration characteristic for Mn-O binding (not present in the diatomite sample). For the Mn/diatomite sample after Pb(II) adsorption, the intensity of the absorption bands is very low, probably because of the adsorbed Pb(II) layer. In particular, no vibrations are observed at wavenumbers lower than 470 cm^-1. This is probably because the Pb(II) ions are bound to the manganese oxide adsorption sites.

Figure 7

FT-IR spectra of natural diatomite, Mn/diatomite, Mn/diatomite after Pb(II) adsorption, and MnO₂: (a) wavenumber region 4000-400 cm^-1; (b) wavenumber region 600-400 cm^-1.

(b)

3.3. Adsorption of Pb(II) on Mn/Diatomite

3.3.1. Adsorption Kinetics of Pb(II) on Mn/Diatomite

Chemical kinetics is indispensable in adsorption studies to determine the adsorption rate of the adsorbate at the solid interface. Kinetic models allow the estimation of adsorption rates and lead to suitable rate expressions and suggestion of possible reaction mechanisms.

We can see that the adsorption efficiency of Pb(II) increases with an increase in contact time and initial concentration (Figure 8). This increase probably depends on the concentration difference of Pb(II) ions on the adsorbent surface and in the solution (driving force). When the initial concentration is small, the driving force is also small, leading to a low adsorption efficiency. The adsorption efficiency increases with the driving force, but this increase becomes smaller when the initial concentration is higher. Figure 8 also shows that the Pb(II) adsorption occurs very rapidly during the first 10 min of contact; then, it decreases slowly in the next 60 min and becomes stable after 240 min. Thus, the time required for the adsorption of Pb(II) on Mn/diatomite to reach the adsorption-desorption equilibrium is 60 min.

Figure 8

Adsorption kinetics of Pb(II) onto Mn/diatomite with the different initial concentrations ( $V = 0.1 L$ ; $m of adsorbent = 0.1 g$ ; temperature: 303 K; $p H_{initial} = 6.19 \div 5.12$ ; $p H_{final} = 6.00 \div 4.53$ ).

The adsorption kinetics were described by using the pseudo-first/second-order kinetic models, and the results are presented in Tables 2 and 3. It was found that the equilibrium adsorption capacity values, $q_{2}$ , calculated from the pseudo-second-order equation, are very close to the experimental equilibrium adsorption capacity values, $q_{e}$ . The kinetic data fit the pseudo-second-order kinetic model better than the pseudo-first-order kinetic model because the former’s RMSE values are smaller, and the data points are closer to the theoretical line than the latter’s (Figure S5). Since the pseudo-second-order kinetic model is derived from the assumption that the rate-limiting step is chemical adsorption, it is concluded that the adsorption process is controlled by chemisorption.

Table 2

Kinetic parameters of Pb(II) adsorption on Mn/diatomite at different initial concentrations of Pb(II) solutions.

Initial Pb(II) concentration (mg L^-1)	$q_{e}$ (mg g^-1)	Pseudo-first-order			Pseudo-second-order
Initial Pb(II) concentration (mg L^-1)	$q_{e}$ (mg g^-1)	$k_{1}$ (min^-1)	$q_{1}$ (mg g^-1)	RMSE	$k_{2}$ (g mg^-1 min^-1)	$q_{2}$ (mg g^-1)	RMSE
31.8	31.5	0.2321	31.1	0.334	0.0272	31.7	0.208
57.6	52.2	0.1085	49.2	1.724	0.0038	52.4	0.689
91.4	74.4	0.1149	68.4	3.113	0.0029	72.8	1.497
115.5	77.9	0.1561	73.1	3.408	0.0041	76.8	1.500
145.1	88.2	0.2037	84.4	2.877	0.0059	87.3	1.373
188.3	90.3	0.1809	86.5	3.150	0.0046	90.0	1.211

Table 3

Pb(II) adsorption kinetic parameters on Mn/diatomite at different adsorption temperatures.

Temperature (K)	$q_{e}$ (mg g^-1)	Pseudo-first-order			Pseudo-second-order
Temperature (K)	$q_{e}$ (mg g^-1)	$k_{1}$ (min^-1)	$q_{1}$ (mg g^-1)	RMSE	$k_{2}$ (g mg^-1 min^-1)	$q_{2}$ (mg g^-1)	RMSE
303	52.4	0.1347	49.8	2.003	0.0052	52.4	0.577
313	61.8	0.1461	58.9	2.810	0.0047	62.1	1.112
323	67.5	0.3386	66.0	0.860	0.0283	66.7	0.517
328	71.8	0.2527	69.7	1.678	0.0115	71.4	0.777
333	74.5	0.3319	72.6	1.412	0.0200	73.7	0.907
338	74.8	0.3345	73.3	0.994	0.0237	74.2	0.536

3.3.2. Effect of Solution pH

Figure 9(a) displays the pH effect on the Pb adsorption efficiency. Overall, it is clear that the adsorption efficiency increases consistently with increasing pH. The adsorption efficiency soars from pH 2 to pH 5 and then increases steadily up to pH 6. The Pb(II) species can be in the forms of Pb²⁺, Pb(OH)⁺, Pb₃(OH)₄ ²⁺, Pb(OH)₂, and Pb(OH)₃ ^-, depending on the solution pH [62]. The cation form Pb²⁺ mainly exists at $pH < 5$ , while hydroxyl anions of lead are predominant at $pH > 5$ . The increasing adsorption at low pH (<pH_PZC, Figure 9(b)) is explained by the fact that the adsorbent surface becomes less positively charged with increasing pH, enhancing the electrostatic interaction between the adsorbate and the surface. At higher pH (>pH_PZC), the adsorption capacity increases slightly because of the electrostatic repulsion of the more negative adsorbate and the negatively charged surface. This fact also indicates that the electrostatic interaction is not the only mechanism of the adsorption process. The pH_PZC of Mn/diatomite after Pb adsorption shifts to low pHs, indicating that the high adsorption efficiency of Mn/diatomite under weak acidic/alkaline conditions could be attributed to the formation of complexes between Pb(II) and the surface hydroxyl groups of the Mn/diatomite. The formation of outer-sphere surface complexes could not shift the point of zero charge of the absorbent because there are no specific chemical reactions among the lead cations that could change the surface charge. The shift of the point of zero charge to a lower pH range indicates the formation of anionic negatively charged surface complexes [63]. Hence, the decrease in pH_PZC to 3.2 implies that the adsorption of lead would be a result of the formation of both outer-sphere complexes and the negatively charged inner-sphere complexes between lead and the adsorbent. Then, the possible adsorption reaction of Mn/diatomite with Pb(II) can be shown schematically as follows:

Figure 9

(a) Pb(II) adsorption capacity on Mn/diatomite at different solution pHs; (b) point of zero charge of Mn/diatomite in 0.01 M KCl solution and in 0.05 mM Pb²⁺ solution determined by the pH drift method [64].

(b)

At low pH, Pb(II) reacts with hydroxyl groups (silanol, Mn-OH) on the surface of Mn/diatomite to form complex adsorption: $\begin{matrix} (13) & Mn / diatomite - OH + {Pb}^{2 +} ⟶ Mn / diatomite - O - {Pb}^{+} + H^{+} \\ 2 (Mn / diatomite - OH) + {Pb}^{2 +} ⟶ {(Mn / diatomite - OH)}_{2} Pb + 2 H^{+} \end{matrix}$

At high pH, hydroxyl anions of lead react with hydroxyl groups (silanol, Mn-OH) on the surface of Mn/diatomite to form complex adsorption: $\begin{matrix} (14) & Mn / diatomite - OH + Pb {(OH)}^{+} ⟶ Mn / diatomite - O - {Pb}^{+} + H_{2} O \\ Mn / diatomite - OH + {Pb}_{3} {(OH)}_{4}^{2 +} ⟶ Mn / diatomite - O - {Pb}_{3} {(OH)}_{3}^{2 +} + H_{2} O \\ Mn / diatomite - OH + Pb {(OH)}_{3^{-}} ⟶ Mn / diatomite - O - Pb {(OH)}_{2^{-}} + H_{2} O \\ 2 (Mn / diatomite - OH) + Pb {(OH)}_{3^{-}} ⟶ {(Mn / diatomite - O)}_{2} Pb {(OH)}^{-} + 2 H_{2} O \end{matrix}$

3.3.3. Adsorption Isotherms

The study of adsorption isotherms is essential in describing the characteristic relationship between the concentration of the adsorbate and the adsorption capacity of the adsorbent, especially when designing an ideal adsorption system in industry [42, 45].

The results of the determination of isotherm parameters of Pb(II) adsorption on Mn/diatomite, natural diatomite, and nano-MnO₂ are presented in Table 4 (and Figure S6). Based on the RMSE value, it can be assumed that the Freundlich and Sips models are more suitable to describe the adsorption of Mn/diatomite and natural diatomite than the Langmuir model. On the contrary, the Sips and Langmuir models are more suitable than the Freundlich model to describe the adsorption on nano-MnO₂. To evaluate the compatibility of models with experimental data, numerous authors have used statistical methods, such as the paired-sample $t$ -test [49], Akaike’s information criteria (AIC) [51], and the sum of normalized errors (SNE) [50]. In general, these methods are based on the relationship between the experimental and calculated adsorption capacities. Figure 10(a) depicts the Langmuir, Freundlich, and Sips isotherms for Pb(II) adsorption on Mn/diatomite. The complete fit occurs when the points lie on the diagonal ( $x = y$ ). It can be seen that the Freundlich and Sips models have a lower dispersion; that is, there are smaller gaps between the prediction of these models and the experimental data. However, the Sips maximum adsorption capacity values ( $q_{m_{S}}$ ) for Mn/diatomite and natural diatomite are much larger than the experimental equilibrium adsorption capacity values (see Tables 2 and 3 and Table S1), indicating that the Sips model does not sufficiently satisfy describing these adsorption processes. Meanwhile, the maximum adsorption capacity values obtained from the Langmuir and Sips equations of nano-MnO₂ are approximately the same, indicating the similarity between these two models.

Table 4

Parameters of the nonlinear models for the adsorption of Pb(II) in an aqueous solution on Mn/diatomite, natural diatomite, and nano-MnO₂.

Isotherm model	Parameter	Adsorbent
Isotherm model	Parameter	Mn/diatomite	Natural diatomite	Nano-MnO₂
Langmuir (2p) $^{*}$	$K_{L}$ (L g^-1)	1.18	4.98	7.54
	$q_{m}$ (mg g^-1)	81.42	16.31	19.42
	RMSE	10.17	1.63	0.73

Freundlich (2p) $^{*}$	$K_{F}$ (mg^(1-1/n) L^1/n g^-1)	41.09	11.20	15.43
	$n$	5.54	5.95	10.35
	RMSE	3.46	0.49	1.45

Sips (3p) $^{*}$	$K_{S}$ (L ^m mg^m-1)	0.27	0.05	18.49
	$q_{m_{S}}$ (mg g^-1)	188.74	240.09	18.93
	$m_{S}$	0.28	0.18	1.58
	RMSE	3.10	0.50	0.62

$^{*}$ Number of parameters.

Figure 10

(a) Relationship between the experimental and the calculated adsorbed amount for the adsorption of Pb(II) on Mn/diatomite; (b) effect of temperature and contact time on the adsorption capacity of Pb(II) in aqueous solution of Mn/diatomite ( $p H_{initial} = 5.48$ ).

(b)

The maximum capacity of monolayer adsorption determined with the nonlinear Langmuir model is 81.42 mg g⁻¹. Our value is slightly smaller than that published by Khraished et al. [11] (99.00 mg g⁻¹) but higher than Li et al.’s value (56.843 mg g⁻¹) [36] and other authors’ values (Table 5). The maximum adsorption capacity of natural diatomite and nano-MnO₂ is 16.31 and 19.42 mg g⁻¹, respectively. These values are significantly smaller than the maximum adsorption capacity of Mn/diatomite. Because a layer of manganese oxide coats the diatomite, this material has an increased adsorption capacity, indicating that Mn/diatomite is not merely a physical mixture of natural diatomite and nano-MnO₂ but is a composite with a high adsorption capacity.

Table 5

Pb(II) adsorption capacity of different adsorbents.

Adsorbent	Adsorption capacity (mg g^–1)	pH	Reference
Mn/diatomite	81.42 (303 K)	Neutral	This work
Natural diatomite	16.31 (303 K)	Neutral	This work
Nano-MnO₂	19.42 (303 K)	Neutral	This work
Mn/diatomite	99 (289 K)	4	[10]
Diatomite	24.94 (296 K)	4	[11]
Mn/diatomite	99 (296 K)	4	[11]
Raw DE	8.5058 (303 K)	Neutral	[36]
MnO₂@DE	56.843 (303 K)	Neutral	[36]
Silicalite-1-NH₂	43.5 (298 K)	5.5	[65]
MCM-48-NH₂	75.2 (298 K)	5.5	[65]
MCM-48-SH	31.2 (298 K)	5.5	[65]
Aminopropyl-MCM-41	64.21 (300 K)	Neutral	[49]
Leonardite	23.89 (303 K)	6	[50]
Biological sludge through ferric activation	42.96 (298 K)	5	[7]
Chitosan-tripolyphosphate bead	57.33 (300 K)	4.5	[44]

3.3.4. Adsorption Thermodynamics

We can see that the adsorption capacity depends on the temperature in a similar manner to that of the initial concentration (Figure 10(b)). At lower temperatures (303–323 K), the maximal capacity increases more rapidly than at higher temperatures (323–338 K), and it remains practically the same at 333 and 338 K. This behaviour is also revealed from the slope of the lines during the first 10 min of adsorption. The higher the temperature, the steeper the line is. For all the temperatures, the adsorption-desorption equilibrium is reached after around 60 min.

The plot of $\ln K_{C}$ against $1 / T$ is shown in Figure 11(a). From the intercept and slope of the regression line, we calculate $Δ H$ ° and $Δ S$ ° from Eq. (10) and $Δ G$ ° from Eq. (11) (Table 6). Figure 11(a) shows that Pb(II) adsorption on Mn/diatomite is strongly dependent on temperature. The temperature range from 303 to 338 K can be divided into two ranges: range 1 from 303 to 323 K and range 2 from 323 to 338 K. In both temperature ranges, $Δ H ° > 0$ , indicating that the adsorption is endothermic; that is, as the temperature increases, the adsorption capacity increases. This result is entirely consistent with the adsorption data shown in Figure 10(b). Furthermore, range 2 has a much larger $Δ H$ ° (234.222 kJ mol⁻¹) than range 1, in which $Δ H ° = 55.451 kJ mo l^{- 1}$ . This suggests that the adsorption process occurs favourably at high temperatures, suggesting chemical adsorption onto the composite.

Figure 11

(a) Linear plot of $\ln (K_{C})$ versus $1 / T$ and (b) Pb(II) adsorption capacity in aqueous solution of reused Mn/diatomite samples (experimental conditions: 0.05 g of the adsorbent and 50 mL of Pb(II) solution with $C_{0} = 80.1 mg L^{- 1}$ ; each experiment was repeated three times).

(b)

Table 6

Thermodynamic parameters of Pb(II) adsorption on Mn/diatomite.

$T$ (K)	$Δ G^{o}$ (kJ mol^-1)	$Δ H^{o}$ (kJ mol^-1)	${Δ S}^{o}$ (J mol^-1 K^-1)
303	-2.111	55.451	189.975
313	-4.011
323	-5.911
323	-5.667	234.222	742.690
328	-9.380
333	-13.094
338	-16.807

The negative value of $Δ G$ ° in the 303–338 K range shows that the Pb(II) adsorption in aqueous solutions on Mn/diatomite is spontaneous. Meanwhile, the positive $Δ S$ ° reflects the increase in the degree of free movement at the solid/solution interface in the adsorption process [13, 43, 44, 46].

Since the Pb(II) adsorption on Mn/diatomite follows the pseudo-second-order kinetic model, the adsorption rate constant $k_{2}$ is used in Eq. (12). Figure S7 depicts the plot of $\ln (k_{2})$ versus $1 / T$ , and the activation energy of Pb(II) adsorption in aqueous solutions on Mn/diatomite material is 41.56 kJ mol⁻¹. Usually, physisorption has an activation energy in the range of 5-40 kJ mol⁻¹, while that of chemisorption is generally in the range of 40-800 kJ mol⁻¹ [46]. Therefore, the adsorption of Pb(II) on Mn/diatomite could occur via the formation of weak chemical bonds between the adsorbent and the adsorbate [41].

3.3.5. Recycling Studies

The regeneration of the spent Mn/diatomite adsorbent was carried out in three adsorption-desorption cycles. The readsorption experiments were performed on 50 mL of Pb(II) solution with pH and initial concentration of 5.45 and 80.1 mg L^-1, respectively, and 0.05 g of the adsorbent. The mixture was shaken with a shaker at 200 rpm for 3 h at ambient temperature ( $29 \pm 2$ °C). Then, the adsorbent was removed by filtration, the remaining Pb(II) concentration was determined, and the adsorption capacity was calculated. The results show that the adsorption capacity of Pb(II) decreases with the number of recycles, and Mn/diatomite could still exhibit approximately 55% of its original adsorption capacity in the 3^rd recycle (Figure 11(b)).

Mn/diatomite after each recycle was subjected to FT-IR measurement, and it was found that the characteristic vibration bands of the diatomite modified by manganese oxide remain (Figure S8a). However, the peaks at the wavenumbers less than 470 cm^-1 are observed only in the first recycled sample (Figure S8b). This indicates that the chemical bonding between Pb(II) and the manganese oxide adsorption sites reduces desorption. This is also the cause of the reduced adsorption capacity of the recycled samples.

4. Conclusions

The manganese-modified diatomite material was successfully prepared with the hydrothermal method from KMnO₄, HCl, and diatomite. The resulting Mn/diatomite possesses a high manganese content (20.47% by weight) and a large specific surface area ( $S_{BET} = 68.5 m^{2} g^{- 1}$ ). The manganese oxide covers the diatomite surface. Mn/diatomite can be used as an adsorbent to remove Pb(II) ions from aqueous solutions with a maximum adsorption capacity of 81.42 mg g⁻¹. The Pb(II) adsorption on Mn/diatomite is spontaneous and endothermic, with an activation energy of 41.56 kJ mol^-1. The material is a stable, potential adsorbent in environmental treatments.

Footnotes

Data Availability

The data used to support the findings of this study are available from the corresponding authors upon request.

Conflicts of Interest

The authors declare that they have no conflicts of interest.

Acknowledgments

The authors would like to thank the Center for Scientific Research and Practice, Institute of Applied Technology, Thu Dau Mot University (Binh Duong province, Vietnam), for the support with the experimental equipment and AAS measurements.

Supplementary Materials

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Removal of Pb(II) from Aqueous Solutions with Manganese Oxide-Modified Diatomite

Abstract

1. Introduction

2. Experimental

2.1. Materials and Chemicals

2.2. Preparation of Mn/Diatomite

2.3. Preparation of Nano-MnO2

2.4. Characterization

2.5. Adsorption Experiments

2.5.1. Evaluate the Pb(II) Adsorption Capacity of the Material

2.5.2. Pb(II) Adsorption Kinetics

2.5.3. Effect of Solution pH

2.5.4. Equilibrium Studies

2.5.5. Thermodynamic Studies

2.5.6. Desorption Experiments

3. Results and Discussion

3.1. Modification of Diatomite with Manganese Oxide and Pb(II) Adsorption Capacity Assessment

3.1.1. Effect of Hydrothermal Temperature

3.1.2. Effect of Molar Ratio of Precursors

3.1.3. Effect of Hydrothermal Time

3.2. Characteristic Properties and Pb(II) Adsorption Capacity of Nano-MnO2

3.3. Adsorption of Pb(II) on Mn/Diatomite

3.3.1. Adsorption Kinetics of Pb(II) on Mn/Diatomite

3.3.2. Effect of Solution pH

3.3.3. Adsorption Isotherms

3.3.4. Adsorption Thermodynamics

3.3.5. Recycling Studies

4. Conclusions

Footnotes

Data Availability

Conflicts of Interest

Acknowledgments

Supplementary Materials

References

2.3. Preparation of Nano-MnO₂

3.2. Characteristic Properties and Pb(II) Adsorption Capacity of Nano-MnO₂