# Synthesis ofα-MnO2 Nanomaterial from a Precursor γ-MnO2: Characterization and Comparative Adsorption of Pb(II) and Fe(III)

Authors: Van-Phuc Dinh; Ngoc-Chung Le; Thi-Phuong-Tu Nguyen; Ngoc-Tuan Nguyen
Journal: Journal of Chemistry
 (2016)

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## Abstract

α-MnO2 nanostructure was successfully synthesized via hydrothermal treatment of a precursor γ-MnO2. Structure, morphology, and BET surface area were characterized using X-ray powder diffraction (XRD), Scanning Electron Microscopy (SEM), and Brunauer-Emmett-Teller nitrogen adsorption (BET-N2 adsorption). Thermal analysis result showed that α-MnO2 nanorods were formed from γ-MnO2 at 600°C. In addition, Pb(II) and Fe(III) adsorptive properties were investigated in an optimal condition. Results showed that equilibrium adsorption was obtained after 60 minutes for Pb(II) at pH = 4.0 and 80 minutes for Fe(III) at pH = 3.5 with 240 rpm of shaking speed overall. Experimental data was analyzed using three models: Langmuir, Freundlich, and Sips. Adsorption capacities (qm) from the Langmuir isotherm models are 124.87 mg/g for Pb(II) and 30.83 mg/g for Fe(III). Along with the highest corelation coefficients, it is clear that the adsorption of Pb(II) and Fe(III) ions on α-MnO2 surface followed Sips model. Kinetic studies indicated that the uptake of Pb(II) and Fe(III) occurred in the pseudo-second-order model with two stages for Pb(II) and three stages for Fe(III).


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## Body

## 1. Introduction

Nowadays, the human health and the environment are in danger owing to the increase in industrial and mining activities [1]. Lead and iron ions seriously affect human life even in trace concentrations. According to the Environmental Quality Act, the limit for lead and iron in water is less than 0.1 mg/L and 1.0 mg/L. Therefore, there are many physicochemical treatments used to remove them, such as chemical precipitation, adsorption and ionic exchange, membrane technology, and solvent extraction. One of the most efficient and promising methods for treatment of lead and iron ions from an aqueous solution is the adsorption technology because of its high enrichment efficiency and ease of phase separation [2].Being well known as a low-cost and environmentally friendly material, manganese oxides have been studied and applied to different areas since they have a great number of crystalline structures (α, β, γ, δ, etc.) and excellent chemical characteristics. Li batteries which are one of the important applications have been widely studied [3–9]. In addition, MnO2 has been used for catalysts [10–13], supercapacitors [7, 14–17], and so forth. In recent years, MnO2 has been studied as an adsorbent to remove inorganic as well as organic compounds from aqueous solution [8, 18–22]. However, adsorptive properties of α-MnO2 nanomaterial as well as comparative adsorption of Pb(II) and Zn(II) onto this material are seldom reported.In this study,α-MnO2 was synthesized from heating precursor γ-MnO2 and used as an adsorbent to remove Pb(II) and Zn(II) from an aqueous solution in an optimal condition. Comparative adsorption of two ions on α-MnO2 was also investigated using three isotherm models: Langmuir, Freundlich, and Sips. The uptake kinetics were analyzed using pseudo-first-order and pseudo-second-order models. In addition, intraparticle diffusion kinetics were also studied to ascertain the mechanism of the uptake.


## 2. Materials and Methods

### 2.1. Preparation ofα-MnO2 from γ-MnO2

γ-MnO2 which was successfully synthesized by Le and Van Phuc [2] from ethanol and potassium permanganate was heated at 600°C. The product after being investigated by XRD, SEM, and BET was used as an absorbent to remove lead and iron ions from the aqueous solution.


### 2.2. Adsorbent

Lead(II) and iron(II) ions were used as adsorbates. 1000 mg/L standard stock solution of each metal ion was prepared by dissolving Pb(NO3)2 and Fe(NO3)3 in distilled water.


### 2.3. Instruments

X-ray diffractometer D5000 made in Germany by Siemens with X-ray radiation CuKα, λ = 1,54056 Å, was used to examine the phase of the crystalline structure.The morphology of the material was investigated by Ultra High Resolution Scanning Electron Microscopy S-4800, a transmission electron microscope.Atomic Absorption Spectrophotometer (Atomic Absorption Spectrometer AA-7000 made in Japan by Shimadzu) was used to determine the BET surface area and pore site.The pH measurements were done with a pH meter (Martini Instruments Mi-150, Romania); the pH meter was standardized using Hanna Instruments buffer solutions with pH values of4.01±0.01, 7.01±0.01, and 10.01±0.01.A temperature-controlled shaker (Model IKA R5) was used for equilibrium studies.


### 2.4. Diagram for These Studies

See Figure1.Figure 1


### 2.5. Adsorption Study

50 mL solution of heavy metal (Pb2+ and Zn2+) ions was placed into a 100 mL conical flask containing 0.1 g α-MnO2. During the uptake, influence of pH of the initial solution was adjusted within 2–5 range using HNO3 0.1 M or NaOH 0.1 M solutions. Adsorption studies were also conducted in batch experiments with various adsorption times (20–240 minutes) and concentrations of metal ion (100–500 mg/L). Concentrations of heavy metal ions in the filtrate before and after uptake were determined using AAS.Adsorption capacity was calculated by using the mass balance equation for the adsorption [1]:(1)q=C0-Ce·Vm.And the adsorption efficiency (%) was calculated from the formula(2)%Removal=C0-Ce·100%C0,where q is the adsorption capacity (mg/g) at equilibrium, C0 and Ce are the initial and the equilibrium concentrations (mg/L), respectively, V is the volume (L) of the solution, and m is the mass (g) of the adsorbent used.


## 2.1. Preparation ofα-MnO2 from γ-MnO2

γ-MnO2 which was successfully synthesized by Le and Van Phuc [2] from ethanol and potassium permanganate was heated at 600°C. The product after being investigated by XRD, SEM, and BET was used as an absorbent to remove lead and iron ions from the aqueous solution.


## 2.2. Adsorbent

Lead(II) and iron(II) ions were used as adsorbates. 1000 mg/L standard stock solution of each metal ion was prepared by dissolving Pb(NO3)2 and Fe(NO3)3 in distilled water.


## 2.3. Instruments

X-ray diffractometer D5000 made in Germany by Siemens with X-ray radiation CuKα, λ = 1,54056 Å, was used to examine the phase of the crystalline structure.The morphology of the material was investigated by Ultra High Resolution Scanning Electron Microscopy S-4800, a transmission electron microscope.Atomic Absorption Spectrophotometer (Atomic Absorption Spectrometer AA-7000 made in Japan by Shimadzu) was used to determine the BET surface area and pore site.The pH measurements were done with a pH meter (Martini Instruments Mi-150, Romania); the pH meter was standardized using Hanna Instruments buffer solutions with pH values of4.01±0.01, 7.01±0.01, and 10.01±0.01.A temperature-controlled shaker (Model IKA R5) was used for equilibrium studies.


## 2.4. Diagram for These Studies

See Figure1.Figure 1


## 2.5. Adsorption Study

50 mL solution of heavy metal (Pb2+ and Zn2+) ions was placed into a 100 mL conical flask containing 0.1 g α-MnO2. During the uptake, influence of pH of the initial solution was adjusted within 2–5 range using HNO3 0.1 M or NaOH 0.1 M solutions. Adsorption studies were also conducted in batch experiments with various adsorption times (20–240 minutes) and concentrations of metal ion (100–500 mg/L). Concentrations of heavy metal ions in the filtrate before and after uptake were determined using AAS.Adsorption capacity was calculated by using the mass balance equation for the adsorption [1]:(1)q=C0-Ce·Vm.And the adsorption efficiency (%) was calculated from the formula(2)%Removal=C0-Ce·100%C0,where q is the adsorption capacity (mg/g) at equilibrium, C0 and Ce are the initial and the equilibrium concentrations (mg/L), respectively, V is the volume (L) of the solution, and m is the mass (g) of the adsorbent used.


## 3. Results and Discussion

### 3.1. Characterization ofα-MnO2 Nanostructures

#### 3.1.1. Thermal Analysis Result

Thermal analysis (Figure2) showed that there were two specific peaks corresponding to the decrease in volume of a product. From 100°C to 300°C, the endothermic process occurred which combined with a significant reduction of volume (approximately 16.40%) showed chemical dehydration as well as melting of material at 186.41°C. From 300°C to 800°C, there was a slight decrease in volume of about 5.40% which was the continued melting. In addition, there was an endothermic peak at 535.93°C corresponding to allotropic transformation. It was predicted that α-MnO2 crystalline structure can be formed at above 536°C.Figure 2
TG-DTA analysis.


#### 3.1.2. X-Ray Analysis Results

XRD analysis results showed thatα-MnO2 crystalline structure was not formed at 400°C (Figure 3). At higher temperatures, 600°C and 800°C, α-MnO2 crystalline structure was completely formed with some specific peaks at 2θ = 28.58, 37.48, 49.78, 59.98, and 68.98 corresponding to the Miller indices (3 1 0), (2 1 1), (4 1 1), (5 2 1), and (5 4 1), respectively (ref. code: 01-072-1982). These results corresponded with prognostication from thermal analysis.Figure 3
XRD image of MnO2 at different temperatures: gamma-MnO2 (a); t = 400°C (b); t = 600°C (c); and t = 800°C (d).


#### 3.1.3. Scanning Electron Microscopy Analysis Results

Figure4 showed that although γ-MnO2 particles were flocculated at 400°C, the morphology of material was not changed. It can be concluded that α-MnO2 nanostructures were not formed. At 600°C and 800°C, the morphology of the material was changed from nanospheres to nanorods which coincided with α-MnO2 crystalline nanostructure corresponding to XRD results. However, these nanorods were broken and smashed at 800°C.Figure 4
SEM images of material at 100°C (a), 400° (b), 600°C ((c), (d)), and 800°C ((e), (f)).
(a) (b) (c) (d) (e) (f)


#### 3.1.4. BET Surface Area and BJH Analysis

The BET surface area as well as BJH adsorption-desorption pore size ofα-MnO2 and γ-MnO2 was provided in Table 1. Whereas precursor γ-MnO2 had a BET surface area of about 65 m2·g−1, the synthesized α-MnO2 had smaller surface, approximately 9.4 m2·g−1. This can be explained by the fact that γ-MnO2 was composed of nanospheres, had a porous surface, and was small in size at room temperature. When heating γ-MnO2 to form α-MnO2 at 600°C, nanospheres were melted, flocculated, and showed an increase in size. As a result, α-MnO2 had a smaller surface area than precursor γ-MnO2. Both γ-MnO2 and α-MnO2 had BJH adsorption average pore width from 20 Å to 50 Å which were mesopores materials. However, with the fact that BJH adsorption pore size of γ-MnO2 is smaller than that of α-MnO2, it was revealed that adsorptive properties of γ-MnO2 were better than those of α-MnO2.Table 1
BET and BJH analysis results ofγ-MnO2 and α-MnO2.
Material BET surface area BJH adsorption pore size BJH desorption pore size γ-MnO2 65.00 m2·g−1 417.83 Å 340.23 Å α-MnO2 9.37 m2·g−1 162.95 Å 734.37 Å


### 3.2. Adsorption Studies

#### 3.2.1. Effective Factors

pH plays an important role in absorbing Pb(II) and Fe(III) ions ontoα-MnO2 material surface. At higher pH value, hydroxo ion MOHn+ or/and insoluble hydroxide MOHn was easily formed which inhibits the adsorption of these ions. In contrast, at low pH value, adsorption sites may be decreased because part of the MnO2 nanomolecules can be dissolved in an acid solution as in the following reaction:(3)MnO2+4H++e-⟶Mn2++2H2OAs a result, adsorption reached a maximum at pH = 4.0 for Pb(II) and pH = 3.5 for Fe(III) (Figure5(a)).Figure 5
Influence of pH (a) and contact time (b) on adsorption of Pb(II) and Fe(III) at room temperature with 240 rpm of shaking speed.
(a) (b)Adsorption time is one of the important factors which helps us to predict kinetics as well as the mechanism of the uptake of heavy metals on material surface. The influence of contact time on the adsorption process of Pb(II) and Fe(III) ontoα-MnO2 was shown in Figure 5(b). As a result, the adsorption equilibrium was obtained after 60 minutes for Pb(II) and 80 minutes for Fe(III).


#### 3.2.2. Isotherm Equation Studies

To understand the nature of the adsorption of Pb(II) and Fe(III) on material surface, some isotherm equations, such as Langmuir, Freundlich, and Sips, were investigated. While Langmuir model can help us to calculate the maximum adsorption capacity on a monolayer, Freundlich model shows the interaction between the absorbent and a multilayer. Plots of these nonlinear equations and equilibrium isotherm parameters were shown in Figure6 and Table 2. Results showed that the maximum adsorption capacities calculated from Langmuir model were 124.87 mg/g for Pb(II) and 30.83 mg/g for Fe(III). This can be explained by the fact that lead ion is more satisfied with adsorption pore size because it has ion radius larger than iron. In addition, the separation factor, SF, which is a dimensionless constant equilibrium parameter, was given by(4)SF=11+KL·C0.Table 2
Equilibrium isotherm parameters.
Isotherm Nonlinear forms Isotherm parameters Pb(II) Fe(III) Langmuir q e = q m · K L · C e 1 + K L · C e K L 0.3340 0.135 q m (mg/g) 124.9 30.83 RMSE 3.529 1.504 R 2 0.9109 0.9042 χ 2 0.6618 0.7603 S F (at the lowest C0) 0.0291 0.0689 S F (at the highest C0) 5.95·10−3 0.0146  Freundlich q e = K F · C e 1 / n 1 / n 0.1115 0.2033 K F 74.76 11.52 RMSE 1.493 0.695 R 2 0.9840 0.9795 χ 2 0.1148 0.1296  Sips q e = K S · C e β S 1 + α S · C e β S K S 6.75 11.84 α S −0.9142 0.1986 β S 0.0070 0.3442 RMSE 1.4689 0.6272 R 2 0.9846 0.9833 χ 2 0.1133 0.1122 q e: adsorption capacity at equilibrium (mg/g); qm: maximum adsorption capacity (mg/g); Ce: equilibrium concentration (mg/L); KL: Langmuir constant; KF: Freundlich constant; n: adsorption intensity; KS: Sips constant; αS: Sips isotherm model constant (L/mg); βS: Sips isotherm model exponent; R2: correlation coefficient; RMSE: root mean square error; χ2: nonlinear chi-square test.Figure 6
Plots of adsorption isotherm models: Langmuir, Freundlich, and Sips models.Based on theSF value, it can be revealed that the smaller the value of SF is, the more favorable the adsorption is. The calculated SF values which were smaller than 1 and higher than 0 in both the lowest and the highest concentrations showed that α-MnO2 was an appropriate adsorbent for Fe(III) and Pb(II). However, SF value of Pb(II) > Fe(III) indicated that Fe(III) will compete for binding sites faster than Pb(II) in a mixed metal ions system.Moreover, the1/n values calculated from Freundlich model for Fe(III) higher than Pb(II) indicated that the interaction between α-MnO2 and Fe(III) is more favorable than α-MnO2 and Pb(II).However, the material surface shows heterogeneity. Hence, the uptake of heavy metal ions onto MnO2 can occur on different mechanisms. Sips model, which is Langmuir and Freundlich models combined, generally described exactly the nature of adsorption. Sips isotherm equation was given by formula [1](5)qe=KS·CeβS1+αS·CeβS,where αS is Sips isotherm model constant (L/mg); βS is Sips isotherm model exponent; KS is Sips constant; Ce (mg/L) is the equilibrium concentration; and qe (mg/g) is adsorption capacity. Sips equation is a three-parameter model, and thus it can be solved by Solver Add-In in Microsoft Excel.In comparison with the correlation coefficients (R2), the root mean square error (RMSE), and the nonlinear chi-square test (χ2) values of the three models, it was clear that the adsorption of Pb(II) and Fe(III) onto α-MnO2 followed Sips model due to the highest R2 value as well as the lowest RMSE and χ2 ones.


#### 3.2.3. Kinetic Studies

Pseudo-first-order and pseudo-second-order models are often used to describe the uptake of heavy metal ions onto material surface for adsorption time. Kinetic parameters give essential information for designing and modeling the adsorption processes. However, the two modes cannot determine clearly the nature of the adsorption of Pb(II) and Fe(III) ontoα-MnO2 nanomaterial. Therefore, intraparticle diffusion model which was developed by Weber and Morris was applied for ascertaining a mechanism of the uptake of Pb(II) and Fe(III) onto α-MnO2. Plots of these kinetic models were shown in Figure 7 and kinetic parameters were given in Table 3.Table 3
Kinetic parameters.
Kinetic models Linear forms Kinetic parameters Pb(II) Fe(III) q e ( e x p ) (mg/g) 109.80 16.80  Pseudo-first-order modeldqdt=K1qe-q log ⁡ q e - q t = log ⁡ q e - K 1 · t 2 ,  303 K 1 (min−1) 0.01428 1.38·10−3 R 2 0.5998 0.6873 q e ( c a l ) (mg/g) 5.433 18.12  Pseudo-second-order modeldqdt=K2qe-q2 t q = 1 K 2 q e 2 + 1 q e · t K 2 (g·mg−1·min−1) 0.01108 4.5·10−3 R 2 1.000 0.9958 q e ( c a l ) (mg/g) 109.89 17.92  Intraparticle diffusion model q t = Kdt1/2 + C K d 1 0.3722 0.0309 K d 2 0.0038 0.1759 K d 3  0.0017Figure 7
Pseudo-first-order kinetic plots for the adsorption of Pb(II) (a) and Fe(III) (c). Pseudo-second-order kinetic plots for the adsorption of Pb(II) (b) and Fe(III) (d). Intraparticle diffusion kinetic plots for the adsorption of Pb(II) (e) and Fe(III) (f).
(a) (b) (c) (d) (e) (f)Results showed that the theoreticalqe values which were calculated from pseudo-second-order kinetic models were more accordant with the experimental values, qe(exp), than pseudo-first-order kinetic ones. In addition, correlation coefficients in these pseudo-second-order kinetic models for both Pb(II) and Fe(III) were higher than 0.9950. It was revealed that pseudo-second-order models were satisfied with describing kinetics of the uptake of Pb(II) and Fe(III) on α-MnO2.Furthermore, intraparticle diffusion models showed that the uptake of Pb(II) followed two stages: firstly, Pb(II) ions were quickly adsorbed onα-MnO2 surface; secondly, there was gradual diffusion of absorbents from surface sites into the inner pores. And it was observed that there are three stages in the sorption of Fe(III) ion on the material surface. In the first one, Fe(III) ion was rapidly transferred from the solution to α-MnO2 surface. In the next one, intraparticle diffusion which is the controlling factor gradually occurred. Lastly, the adsorption equilibrium was obtained to correspond with the saturation of absorbent sites. The intraparticle diffusion constants for all these stages were given in Table 3.


## 3.1. Characterization ofα-MnO2 Nanostructures

### 3.1.1. Thermal Analysis Result

Thermal analysis (Figure2) showed that there were two specific peaks corresponding to the decrease in volume of a product. From 100°C to 300°C, the endothermic process occurred which combined with a significant reduction of volume (approximately 16.40%) showed chemical dehydration as well as melting of material at 186.41°C. From 300°C to 800°C, there was a slight decrease in volume of about 5.40% which was the continued melting. In addition, there was an endothermic peak at 535.93°C corresponding to allotropic transformation. It was predicted that α-MnO2 crystalline structure can be formed at above 536°C.Figure 2
TG-DTA analysis.


### 3.1.2. X-Ray Analysis Results

XRD analysis results showed thatα-MnO2 crystalline structure was not formed at 400°C (Figure 3). At higher temperatures, 600°C and 800°C, α-MnO2 crystalline structure was completely formed with some specific peaks at 2θ = 28.58, 37.48, 49.78, 59.98, and 68.98 corresponding to the Miller indices (3 1 0), (2 1 1), (4 1 1), (5 2 1), and (5 4 1), respectively (ref. code: 01-072-1982). These results corresponded with prognostication from thermal analysis.Figure 3
XRD image of MnO2 at different temperatures: gamma-MnO2 (a); t = 400°C (b); t = 600°C (c); and t = 800°C (d).


### 3.1.3. Scanning Electron Microscopy Analysis Results

Figure4 showed that although γ-MnO2 particles were flocculated at 400°C, the morphology of material was not changed. It can be concluded that α-MnO2 nanostructures were not formed. At 600°C and 800°C, the morphology of the material was changed from nanospheres to nanorods which coincided with α-MnO2 crystalline nanostructure corresponding to XRD results. However, these nanorods were broken and smashed at 800°C.Figure 4
SEM images of material at 100°C (a), 400° (b), 600°C ((c), (d)), and 800°C ((e), (f)).
(a) (b) (c) (d) (e) (f)


### 3.1.4. BET Surface Area and BJH Analysis

The BET surface area as well as BJH adsorption-desorption pore size ofα-MnO2 and γ-MnO2 was provided in Table 1. Whereas precursor γ-MnO2 had a BET surface area of about 65 m2·g−1, the synthesized α-MnO2 had smaller surface, approximately 9.4 m2·g−1. This can be explained by the fact that γ-MnO2 was composed of nanospheres, had a porous surface, and was small in size at room temperature. When heating γ-MnO2 to form α-MnO2 at 600°C, nanospheres were melted, flocculated, and showed an increase in size. As a result, α-MnO2 had a smaller surface area than precursor γ-MnO2. Both γ-MnO2 and α-MnO2 had BJH adsorption average pore width from 20 Å to 50 Å which were mesopores materials. However, with the fact that BJH adsorption pore size of γ-MnO2 is smaller than that of α-MnO2, it was revealed that adsorptive properties of γ-MnO2 were better than those of α-MnO2.Table 1
BET and BJH analysis results ofγ-MnO2 and α-MnO2.
Material BET surface area BJH adsorption pore size BJH desorption pore size γ-MnO2 65.00 m2·g−1 417.83 Å 340.23 Å α-MnO2 9.37 m2·g−1 162.95 Å 734.37 Å


## 3.1.1. Thermal Analysis Result

Thermal analysis (Figure2) showed that there were two specific peaks corresponding to the decrease in volume of a product. From 100°C to 300°C, the endothermic process occurred which combined with a significant reduction of volume (approximately 16.40%) showed chemical dehydration as well as melting of material at 186.41°C. From 300°C to 800°C, there was a slight decrease in volume of about 5.40% which was the continued melting. In addition, there was an endothermic peak at 535.93°C corresponding to allotropic transformation. It was predicted that α-MnO2 crystalline structure can be formed at above 536°C.Figure 2
TG-DTA analysis.


## 3.1.2. X-Ray Analysis Results

XRD analysis results showed thatα-MnO2 crystalline structure was not formed at 400°C (Figure 3). At higher temperatures, 600°C and 800°C, α-MnO2 crystalline structure was completely formed with some specific peaks at 2θ = 28.58, 37.48, 49.78, 59.98, and 68.98 corresponding to the Miller indices (3 1 0), (2 1 1), (4 1 1), (5 2 1), and (5 4 1), respectively (ref. code: 01-072-1982). These results corresponded with prognostication from thermal analysis.Figure 3
XRD image of MnO2 at different temperatures: gamma-MnO2 (a); t = 400°C (b); t = 600°C (c); and t = 800°C (d).


## 3.1.3. Scanning Electron Microscopy Analysis Results

Figure4 showed that although γ-MnO2 particles were flocculated at 400°C, the morphology of material was not changed. It can be concluded that α-MnO2 nanostructures were not formed. At 600°C and 800°C, the morphology of the material was changed from nanospheres to nanorods which coincided with α-MnO2 crystalline nanostructure corresponding to XRD results. However, these nanorods were broken and smashed at 800°C.Figure 4
SEM images of material at 100°C (a), 400° (b), 600°C ((c), (d)), and 800°C ((e), (f)).
(a) (b) (c) (d) (e) (f)


## 3.1.4. BET Surface Area and BJH Analysis

The BET surface area as well as BJH adsorption-desorption pore size ofα-MnO2 and γ-MnO2 was provided in Table 1. Whereas precursor γ-MnO2 had a BET surface area of about 65 m2·g−1, the synthesized α-MnO2 had smaller surface, approximately 9.4 m2·g−1. This can be explained by the fact that γ-MnO2 was composed of nanospheres, had a porous surface, and was small in size at room temperature. When heating γ-MnO2 to form α-MnO2 at 600°C, nanospheres were melted, flocculated, and showed an increase in size. As a result, α-MnO2 had a smaller surface area than precursor γ-MnO2. Both γ-MnO2 and α-MnO2 had BJH adsorption average pore width from 20 Å to 50 Å which were mesopores materials. However, with the fact that BJH adsorption pore size of γ-MnO2 is smaller than that of α-MnO2, it was revealed that adsorptive properties of γ-MnO2 were better than those of α-MnO2.Table 1
BET and BJH analysis results ofγ-MnO2 and α-MnO2.
Material BET surface area BJH adsorption pore size BJH desorption pore size γ-MnO2 65.00 m2·g−1 417.83 Å 340.23 Å α-MnO2 9.37 m2·g−1 162.95 Å 734.37 Å


## 3.2. Adsorption Studies

### 3.2.1. Effective Factors

pH plays an important role in absorbing Pb(II) and Fe(III) ions ontoα-MnO2 material surface. At higher pH value, hydroxo ion MOHn+ or/and insoluble hydroxide MOHn was easily formed which inhibits the adsorption of these ions. In contrast, at low pH value, adsorption sites may be decreased because part of the MnO2 nanomolecules can be dissolved in an acid solution as in the following reaction:(3)MnO2+4H++e-⟶Mn2++2H2OAs a result, adsorption reached a maximum at pH = 4.0 for Pb(II) and pH = 3.5 for Fe(III) (Figure5(a)).Figure 5
Influence of pH (a) and contact time (b) on adsorption of Pb(II) and Fe(III) at room temperature with 240 rpm of shaking speed.
(a) (b)Adsorption time is one of the important factors which helps us to predict kinetics as well as the mechanism of the uptake of heavy metals on material surface. The influence of contact time on the adsorption process of Pb(II) and Fe(III) ontoα-MnO2 was shown in Figure 5(b). As a result, the adsorption equilibrium was obtained after 60 minutes for Pb(II) and 80 minutes for Fe(III).


### 3.2.2. Isotherm Equation Studies

To understand the nature of the adsorption of Pb(II) and Fe(III) on material surface, some isotherm equations, such as Langmuir, Freundlich, and Sips, were investigated. While Langmuir model can help us to calculate the maximum adsorption capacity on a monolayer, Freundlich model shows the interaction between the absorbent and a multilayer. Plots of these nonlinear equations and equilibrium isotherm parameters were shown in Figure6 and Table 2. Results showed that the maximum adsorption capacities calculated from Langmuir model were 124.87 mg/g for Pb(II) and 30.83 mg/g for Fe(III). This can be explained by the fact that lead ion is more satisfied with adsorption pore size because it has ion radius larger than iron. In addition, the separation factor, SF, which is a dimensionless constant equilibrium parameter, was given by(4)SF=11+KL·C0.Table 2
Equilibrium isotherm parameters.
Isotherm Nonlinear forms Isotherm parameters Pb(II) Fe(III) Langmuir q e = q m · K L · C e 1 + K L · C e K L 0.3340 0.135 q m (mg/g) 124.9 30.83 RMSE 3.529 1.504 R 2 0.9109 0.9042 χ 2 0.6618 0.7603 S F (at the lowest C0) 0.0291 0.0689 S F (at the highest C0) 5.95·10−3 0.0146  Freundlich q e = K F · C e 1 / n 1 / n 0.1115 0.2033 K F 74.76 11.52 RMSE 1.493 0.695 R 2 0.9840 0.9795 χ 2 0.1148 0.1296  Sips q e = K S · C e β S 1 + α S · C e β S K S 6.75 11.84 α S −0.9142 0.1986 β S 0.0070 0.3442 RMSE 1.4689 0.6272 R 2 0.9846 0.9833 χ 2 0.1133 0.1122 q e: adsorption capacity at equilibrium (mg/g); qm: maximum adsorption capacity (mg/g); Ce: equilibrium concentration (mg/L); KL: Langmuir constant; KF: Freundlich constant; n: adsorption intensity; KS: Sips constant; αS: Sips isotherm model constant (L/mg); βS: Sips isotherm model exponent; R2: correlation coefficient; RMSE: root mean square error; χ2: nonlinear chi-square test.Figure 6
Plots of adsorption isotherm models: Langmuir, Freundlich, and Sips models.Based on theSF value, it can be revealed that the smaller the value of SF is, the more favorable the adsorption is. The calculated SF values which were smaller than 1 and higher than 0 in both the lowest and the highest concentrations showed that α-MnO2 was an appropriate adsorbent for Fe(III) and Pb(II). However, SF value of Pb(II) > Fe(III) indicated that Fe(III) will compete for binding sites faster than Pb(II) in a mixed metal ions system.Moreover, the1/n values calculated from Freundlich model for Fe(III) higher than Pb(II) indicated that the interaction between α-MnO2 and Fe(III) is more favorable than α-MnO2 and Pb(II).However, the material surface shows heterogeneity. Hence, the uptake of heavy metal ions onto MnO2 can occur on different mechanisms. Sips model, which is Langmuir and Freundlich models combined, generally described exactly the nature of adsorption. Sips isotherm equation was given by formula [1](5)qe=KS·CeβS1+αS·CeβS,where αS is Sips isotherm model constant (L/mg); βS is Sips isotherm model exponent; KS is Sips constant; Ce (mg/L) is the equilibrium concentration; and qe (mg/g) is adsorption capacity. Sips equation is a three-parameter model, and thus it can be solved by Solver Add-In in Microsoft Excel.In comparison with the correlation coefficients (R2), the root mean square error (RMSE), and the nonlinear chi-square test (χ2) values of the three models, it was clear that the adsorption of Pb(II) and Fe(III) onto α-MnO2 followed Sips model due to the highest R2 value as well as the lowest RMSE and χ2 ones.


### 3.2.3. Kinetic Studies

Pseudo-first-order and pseudo-second-order models are often used to describe the uptake of heavy metal ions onto material surface for adsorption time. Kinetic parameters give essential information for designing and modeling the adsorption processes. However, the two modes cannot determine clearly the nature of the adsorption of Pb(II) and Fe(III) ontoα-MnO2 nanomaterial. Therefore, intraparticle diffusion model which was developed by Weber and Morris was applied for ascertaining a mechanism of the uptake of Pb(II) and Fe(III) onto α-MnO2. Plots of these kinetic models were shown in Figure 7 and kinetic parameters were given in Table 3.Table 3
Kinetic parameters.
Kinetic models Linear forms Kinetic parameters Pb(II) Fe(III) q e ( e x p ) (mg/g) 109.80 16.80  Pseudo-first-order modeldqdt=K1qe-q log ⁡ q e - q t = log ⁡ q e - K 1 · t 2 ,  303 K 1 (min−1) 0.01428 1.38·10−3 R 2 0.5998 0.6873 q e ( c a l ) (mg/g) 5.433 18.12  Pseudo-second-order modeldqdt=K2qe-q2 t q = 1 K 2 q e 2 + 1 q e · t K 2 (g·mg−1·min−1) 0.01108 4.5·10−3 R 2 1.000 0.9958 q e ( c a l ) (mg/g) 109.89 17.92  Intraparticle diffusion model q t = Kdt1/2 + C K d 1 0.3722 0.0309 K d 2 0.0038 0.1759 K d 3  0.0017Figure 7
Pseudo-first-order kinetic plots for the adsorption of Pb(II) (a) and Fe(III) (c). Pseudo-second-order kinetic plots for the adsorption of Pb(II) (b) and Fe(III) (d). Intraparticle diffusion kinetic plots for the adsorption of Pb(II) (e) and Fe(III) (f).
(a) (b) (c) (d) (e) (f)Results showed that the theoreticalqe values which were calculated from pseudo-second-order kinetic models were more accordant with the experimental values, qe(exp), than pseudo-first-order kinetic ones. In addition, correlation coefficients in these pseudo-second-order kinetic models for both Pb(II) and Fe(III) were higher than 0.9950. It was revealed that pseudo-second-order models were satisfied with describing kinetics of the uptake of Pb(II) and Fe(III) on α-MnO2.Furthermore, intraparticle diffusion models showed that the uptake of Pb(II) followed two stages: firstly, Pb(II) ions were quickly adsorbed onα-MnO2 surface; secondly, there was gradual diffusion of absorbents from surface sites into the inner pores. And it was observed that there are three stages in the sorption of Fe(III) ion on the material surface. In the first one, Fe(III) ion was rapidly transferred from the solution to α-MnO2 surface. In the next one, intraparticle diffusion which is the controlling factor gradually occurred. Lastly, the adsorption equilibrium was obtained to correspond with the saturation of absorbent sites. The intraparticle diffusion constants for all these stages were given in Table 3.


## 3.2.1. Effective Factors

pH plays an important role in absorbing Pb(II) and Fe(III) ions ontoα-MnO2 material surface. At higher pH value, hydroxo ion MOHn+ or/and insoluble hydroxide MOHn was easily formed which inhibits the adsorption of these ions. In contrast, at low pH value, adsorption sites may be decreased because part of the MnO2 nanomolecules can be dissolved in an acid solution as in the following reaction:(3)MnO2+4H++e-⟶Mn2++2H2OAs a result, adsorption reached a maximum at pH = 4.0 for Pb(II) and pH = 3.5 for Fe(III) (Figure5(a)).Figure 5
Influence of pH (a) and contact time (b) on adsorption of Pb(II) and Fe(III) at room temperature with 240 rpm of shaking speed.
(a) (b)Adsorption time is one of the important factors which helps us to predict kinetics as well as the mechanism of the uptake of heavy metals on material surface. The influence of contact time on the adsorption process of Pb(II) and Fe(III) ontoα-MnO2 was shown in Figure 5(b). As a result, the adsorption equilibrium was obtained after 60 minutes for Pb(II) and 80 minutes for Fe(III).


## 3.2.2. Isotherm Equation Studies

To understand the nature of the adsorption of Pb(II) and Fe(III) on material surface, some isotherm equations, such as Langmuir, Freundlich, and Sips, were investigated. While Langmuir model can help us to calculate the maximum adsorption capacity on a monolayer, Freundlich model shows the interaction between the absorbent and a multilayer. Plots of these nonlinear equations and equilibrium isotherm parameters were shown in Figure6 and Table 2. Results showed that the maximum adsorption capacities calculated from Langmuir model were 124.87 mg/g for Pb(II) and 30.83 mg/g for Fe(III). This can be explained by the fact that lead ion is more satisfied with adsorption pore size because it has ion radius larger than iron. In addition, the separation factor, SF, which is a dimensionless constant equilibrium parameter, was given by(4)SF=11+KL·C0.Table 2
Equilibrium isotherm parameters.
Isotherm Nonlinear forms Isotherm parameters Pb(II) Fe(III) Langmuir q e = q m · K L · C e 1 + K L · C e K L 0.3340 0.135 q m (mg/g) 124.9 30.83 RMSE 3.529 1.504 R 2 0.9109 0.9042 χ 2 0.6618 0.7603 S F (at the lowest C0) 0.0291 0.0689 S F (at the highest C0) 5.95·10−3 0.0146  Freundlich q e = K F · C e 1 / n 1 / n 0.1115 0.2033 K F 74.76 11.52 RMSE 1.493 0.695 R 2 0.9840 0.9795 χ 2 0.1148 0.1296  Sips q e = K S · C e β S 1 + α S · C e β S K S 6.75 11.84 α S −0.9142 0.1986 β S 0.0070 0.3442 RMSE 1.4689 0.6272 R 2 0.9846 0.9833 χ 2 0.1133 0.1122 q e: adsorption capacity at equilibrium (mg/g); qm: maximum adsorption capacity (mg/g); Ce: equilibrium concentration (mg/L); KL: Langmuir constant; KF: Freundlich constant; n: adsorption intensity; KS: Sips constant; αS: Sips isotherm model constant (L/mg); βS: Sips isotherm model exponent; R2: correlation coefficient; RMSE: root mean square error; χ2: nonlinear chi-square test.Figure 6
Plots of adsorption isotherm models: Langmuir, Freundlich, and Sips models.Based on theSF value, it can be revealed that the smaller the value of SF is, the more favorable the adsorption is. The calculated SF values which were smaller than 1 and higher than 0 in both the lowest and the highest concentrations showed that α-MnO2 was an appropriate adsorbent for Fe(III) and Pb(II). However, SF value of Pb(II) > Fe(III) indicated that Fe(III) will compete for binding sites faster than Pb(II) in a mixed metal ions system.Moreover, the1/n values calculated from Freundlich model for Fe(III) higher than Pb(II) indicated that the interaction between α-MnO2 and Fe(III) is more favorable than α-MnO2 and Pb(II).However, the material surface shows heterogeneity. Hence, the uptake of heavy metal ions onto MnO2 can occur on different mechanisms. Sips model, which is Langmuir and Freundlich models combined, generally described exactly the nature of adsorption. Sips isotherm equation was given by formula [1](5)qe=KS·CeβS1+αS·CeβS,where αS is Sips isotherm model constant (L/mg); βS is Sips isotherm model exponent; KS is Sips constant; Ce (mg/L) is the equilibrium concentration; and qe (mg/g) is adsorption capacity. Sips equation is a three-parameter model, and thus it can be solved by Solver Add-In in Microsoft Excel.In comparison with the correlation coefficients (R2), the root mean square error (RMSE), and the nonlinear chi-square test (χ2) values of the three models, it was clear that the adsorption of Pb(II) and Fe(III) onto α-MnO2 followed Sips model due to the highest R2 value as well as the lowest RMSE and χ2 ones.


## 3.2.3. Kinetic Studies

Pseudo-first-order and pseudo-second-order models are often used to describe the uptake of heavy metal ions onto material surface for adsorption time. Kinetic parameters give essential information for designing and modeling the adsorption processes. However, the two modes cannot determine clearly the nature of the adsorption of Pb(II) and Fe(III) ontoα-MnO2 nanomaterial. Therefore, intraparticle diffusion model which was developed by Weber and Morris was applied for ascertaining a mechanism of the uptake of Pb(II) and Fe(III) onto α-MnO2. Plots of these kinetic models were shown in Figure 7 and kinetic parameters were given in Table 3.Table 3
Kinetic parameters.
Kinetic models Linear forms Kinetic parameters Pb(II) Fe(III) q e ( e x p ) (mg/g) 109.80 16.80  Pseudo-first-order modeldqdt=K1qe-q log ⁡ q e - q t = log ⁡ q e - K 1 · t 2 ,  303 K 1 (min−1) 0.01428 1.38·10−3 R 2 0.5998 0.6873 q e ( c a l ) (mg/g) 5.433 18.12  Pseudo-second-order modeldqdt=K2qe-q2 t q = 1 K 2 q e 2 + 1 q e · t K 2 (g·mg−1·min−1) 0.01108 4.5·10−3 R 2 1.000 0.9958 q e ( c a l ) (mg/g) 109.89 17.92  Intraparticle diffusion model q t = Kdt1/2 + C K d 1 0.3722 0.0309 K d 2 0.0038 0.1759 K d 3  0.0017Figure 7
Pseudo-first-order kinetic plots for the adsorption of Pb(II) (a) and Fe(III) (c). Pseudo-second-order kinetic plots for the adsorption of Pb(II) (b) and Fe(III) (d). Intraparticle diffusion kinetic plots for the adsorption of Pb(II) (e) and Fe(III) (f).
(a) (b) (c) (d) (e) (f)Results showed that the theoreticalqe values which were calculated from pseudo-second-order kinetic models were more accordant with the experimental values, qe(exp), than pseudo-first-order kinetic ones. In addition, correlation coefficients in these pseudo-second-order kinetic models for both Pb(II) and Fe(III) were higher than 0.9950. It was revealed that pseudo-second-order models were satisfied with describing kinetics of the uptake of Pb(II) and Fe(III) on α-MnO2.Furthermore, intraparticle diffusion models showed that the uptake of Pb(II) followed two stages: firstly, Pb(II) ions were quickly adsorbed onα-MnO2 surface; secondly, there was gradual diffusion of absorbents from surface sites into the inner pores. And it was observed that there are three stages in the sorption of Fe(III) ion on the material surface. In the first one, Fe(III) ion was rapidly transferred from the solution to α-MnO2 surface. In the next one, intraparticle diffusion which is the controlling factor gradually occurred. Lastly, the adsorption equilibrium was obtained to correspond with the saturation of absorbent sites. The intraparticle diffusion constants for all these stages were given in Table 3.


## 4. Conclusion

In this report,α-MnO2 nanostructure synthesized from a precursor γ-MnO2 was used as an adsorbent to remove Pb(II) and Fe(III) from an aqueous solution at an optimal condition. Results showed that the maximum capacities were 124.87 mg/g for Pb(II) and 30.83 mg/g for Fe(III) at pH = 4.0 for Pb(II) and pH = 3.5 for Fe(III). Experiment data which was analyzed using three isotherm models, that is, Langmuir, Freundlich, and Sips, indicated that Sips model was fitted better than the other two. Kinetics studies showed that both Pb(II) and Fe(III) corresponded with pseudo-second-order model with corelation coefficient constants (R2) higher than 0.9950. Intraparticle diffusion kinetics confirmed that the uptake of Pb(II) on α-MnO2 includes two stages while it includes three stages for Fe(III).



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