# Synthesis of MnO2 Hollow Nanospheres through Selective Etching Method as an Effective Absorbent to Remove Methyl Orange from Aqueous Solution

Authors: Xiaohui Hu; Fangsong Zhang; Hong Wang; Xinyi Zhang; Linfeng Zhang
Journal: Journal of Nanomaterials
 (2019)

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## Abstract

The discharge of dye wastewater has become an unavoidable problem for human health and the environment. Developing an economical and rapid method to prepare effective adsorbents for selective removal of dyes is extremely urgent. In this work, MnO2 hollow nanospheres (MHNSs) were prepared through the selective etching method with the MnCO3 as the sacrificial template. The effect of the pH value, contact time, and initial concentration on the adsorption of methyl orange (MO) onto the MHNSs was systematically investigated. The unique mesoporous hollow structure and large BET surface area (43.74 m2/g) of MHNSs lead to an excellent adsorption capacity (1677.14 mg/g) at the optimal condition. Furthermore, the prepared MHNSs also showed great stability (90% removal rate after four cycles). The adsorption kinetics data fitted well with the pseudo-second-order kinetic model (R2>0.9997). The overall process was jointly controlled by external mass transfer and intraparticle diffusion, and intraparticle diffusion was the dominant factor. The adsorption isotherm results showed that the Freundlich model was more accurate to describe the experimental data than the Langmuir model. The thermodynamic analysis showed that the adsorption of MO on MHNSs was spontaneous and exothermic. Moreover, the calculated ΔG° and the XPS spectra showed that the process was mainly a physical process. It is expected that MHNS has a potential application for purifying dye wastewater due to its great adsorption performance and excellent stability.


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## Body

## 1. Introduction

Water pollution has become a worrying challenge for human health and the ecosystem [1–7]. The release of dye wastewater leads to serious environmental problems. However, the treatment of this class of wastewater has proven to be quite difficult because of its stable and complex aromatic structure [8–12]. Therefore, there is an urgent need to design an efficient method for removing dye molecules from wastewater. Various technologies such as photoelectrocatalysis, membrane filtration, chemical oxidation, and adsorption have been proposed to solve this problem [13–18]. Nevertheless, these methods are often expensive, technically demanding, or only effective at a certain concentration. The adsorption technology has been considered to be an efficient, nontoxic, and economic water treatment technique, and the key of this technology is to develop highly effective adsorbents.Currently, as one of typical metal oxides, manganese oxide with various structures and morphologies was widely used in the fields of catalysis, electrochemical supercapacitors, and adsorption. The morphology and structure of prepared manganese oxide are seriously restricted by its synthetic route and conditions. Many researchers have prepared manganese oxide composite adsorbents with uniform microstructure and highly specific surface area through molten salt, sol-gel, coprecipitation, and hydrothermal processes [19–23]. Among reported materials, the hollow micro-/nanostructures are a special class of materials due to their unique morphologies, which have superior advantages such as high BET surface area, surface permeability, and potential scale-dependence [24–27]. Gheju et al. have reported an efficient MnO2 adsorbent for adsorption of Cr6+ and investigated the influence of the pH value, temperature, and initial Cr6+ concentration [28]. The highest adsorption capacity was obtained at 20°C and pH 5.9. Li. et al. have successfully synthesized hierarchical hollow MnO2 microspheres by a hydrothermal method, and the prepared materials show high performance for the adsorption of benzene [29]. However, conventional MnO2-based adsorbents were usually irregular nanoparticles or required a high temperature/pressure condition. Therefore, it is necessary to design a simple operation and short-reaction-time method for preparing a MnO2 adsorbent with high efficiency.In this study, hollow manganese dioxide adsorbents with high efficiency were prepared by a low-temperature acid etching method [30]. The size and shell thickness of the hollow manganese oxide could be exactly controlled by adjusting the amount of etch acid and reaction time. Subsequently, the adsorption capacity of prepared hollow manganese oxide microspheres under different pHs, temperatures, and initial concentrations was evaluated by the adsorption of MO. The adsorption kinetics, rate-determining step, and thermodynamic analysis of the adsorption process were also investigated. Moreover, the possible mechanism of a MO molecule absorbed onto the MHNSs was studied and discussed.


## 2. Experimental

### 2.1. Materials

Methyl orange, sodium hydroxide, hydrochloric acid, potassium permanganate, manganese nitrate, ammonium carbonate, and polyethylene glycol were all purchased from Sinopharm Chemical Reagent Co. Ltd.


### 2.2. Synthesis of the MHNSs

Typically, 1.6 g Mn(NO3)2 (50% solution) and 0.5 g PEG powder were dissolved in 100 mL of ethanol-water solution (v:v=1:1) under 40°C water bath. Afterward, 1.6 g (NH4)2CO3 powder was dissolved into 20 mL deionized water and then transferred to a 20 mL syringe. The above (NH4)2CO3 solution was slowly dripped into the manganese nitrate solution, and the entire drip process was completed within 2 h. After aging for 4 h, the product was collected, washed, and fully dried at 60°C. Then, 0.5 g of prepared MnCO3 powder was dispersed in 20 mL of distilled water and 25 mL of 0.03 M KMnO4 solution was added into the suspension. Then, 8 mL 2.5 M hydrochloric acid solution was added into the mixture with fierce stirring. The etching reaction was sustained for 2 min, then the product was filtered immediately and washed with distilled water. Finally, the obtained solid was dried in a vacuum oven at 60°C for 12 h.


### 2.3. Characterization of the Samples

The crystal structure of the prepared MnCO3 template and MHNSs was studied with X-ray powder diffraction (XRD, Bruker D8 Advance, Cu Kα, λ=0.15418nm). The morphology of the samples was characterized by field emission scanning electron microscopy (FESEM, ZEISS Merlin compact) and transmission electron microscopy (TEM, TecnaiG2 F2o S-TWIN). The surface functional groups were analyzed by Fourier transform infrared (FT-IR) spectra (Nicolet 5700). The BET surface area, pore volume, and pore size distribution of the MHNSs characterized at 77 K were investigated using an Autosorb-iQ Quantachrome surface area and porosity analyzer. The surface compositions were analyzed by X-ray photoelectron spectroscopy (XPS, ESCALab 250Xi).


### 2.4. Evaluation of the Adsorption Capacity

The MO solution was used as a model pollutant to evaluate the adsorption performance of the as-prepared MHNSs. In order to study the effect of initial concentrations, the MO solution with selected concentrations (50, 100, 150, 200, 250, and 300 mg L-1) was prepared. Similarly, the pH value of the solution was adjusted from 2.0 to 11.0 using 1 M HCl or NaOH to investigate the effect of pH value. The 20 mg MHNS adsorbent was added into 200 mL MO solution, and the solution was kept at 15°C under stirring at 300 rpm. At selected intervals, about 3 mL suspension was collected and filtered to remove the catalysts using a 0.45 μm membrane filtration. The MO concentration was recorded on a Lambda 35 UV-Vis spectrometer. The equilibrium absorption capacity of the MHNSs qe (mg g-1) was calculated using the following equation [31, 32]:
(1)qe=VC0−Cem,where C0 (mg/L) is the initial MO concentration, Ce (mg/L) is the equilibrium MO concentration, V (L) is the volume of the solution, and m (g) is the mass of the added MHNS adsorbents.The reusability of the MHNSs was a very important factor for industrial application. In the stability evaluation, the used adsorbent was collected. Then, the stability test of the MHNSs was conducted by repeating the MO solution under the same conditions for five cycles without any regeneration treatment [33, 34].


## 2.1. Materials

Methyl orange, sodium hydroxide, hydrochloric acid, potassium permanganate, manganese nitrate, ammonium carbonate, and polyethylene glycol were all purchased from Sinopharm Chemical Reagent Co. Ltd.


## 2.2. Synthesis of the MHNSs

Typically, 1.6 g Mn(NO3)2 (50% solution) and 0.5 g PEG powder were dissolved in 100 mL of ethanol-water solution (v:v=1:1) under 40°C water bath. Afterward, 1.6 g (NH4)2CO3 powder was dissolved into 20 mL deionized water and then transferred to a 20 mL syringe. The above (NH4)2CO3 solution was slowly dripped into the manganese nitrate solution, and the entire drip process was completed within 2 h. After aging for 4 h, the product was collected, washed, and fully dried at 60°C. Then, 0.5 g of prepared MnCO3 powder was dispersed in 20 mL of distilled water and 25 mL of 0.03 M KMnO4 solution was added into the suspension. Then, 8 mL 2.5 M hydrochloric acid solution was added into the mixture with fierce stirring. The etching reaction was sustained for 2 min, then the product was filtered immediately and washed with distilled water. Finally, the obtained solid was dried in a vacuum oven at 60°C for 12 h.


## 2.3. Characterization of the Samples

The crystal structure of the prepared MnCO3 template and MHNSs was studied with X-ray powder diffraction (XRD, Bruker D8 Advance, Cu Kα, λ=0.15418nm). The morphology of the samples was characterized by field emission scanning electron microscopy (FESEM, ZEISS Merlin compact) and transmission electron microscopy (TEM, TecnaiG2 F2o S-TWIN). The surface functional groups were analyzed by Fourier transform infrared (FT-IR) spectra (Nicolet 5700). The BET surface area, pore volume, and pore size distribution of the MHNSs characterized at 77 K were investigated using an Autosorb-iQ Quantachrome surface area and porosity analyzer. The surface compositions were analyzed by X-ray photoelectron spectroscopy (XPS, ESCALab 250Xi).


## 2.4. Evaluation of the Adsorption Capacity

The MO solution was used as a model pollutant to evaluate the adsorption performance of the as-prepared MHNSs. In order to study the effect of initial concentrations, the MO solution with selected concentrations (50, 100, 150, 200, 250, and 300 mg L-1) was prepared. Similarly, the pH value of the solution was adjusted from 2.0 to 11.0 using 1 M HCl or NaOH to investigate the effect of pH value. The 20 mg MHNS adsorbent was added into 200 mL MO solution, and the solution was kept at 15°C under stirring at 300 rpm. At selected intervals, about 3 mL suspension was collected and filtered to remove the catalysts using a 0.45 μm membrane filtration. The MO concentration was recorded on a Lambda 35 UV-Vis spectrometer. The equilibrium absorption capacity of the MHNSs qe (mg g-1) was calculated using the following equation [31, 32]:
(1)qe=VC0−Cem,where C0 (mg/L) is the initial MO concentration, Ce (mg/L) is the equilibrium MO concentration, V (L) is the volume of the solution, and m (g) is the mass of the added MHNS adsorbents.The reusability of the MHNSs was a very important factor for industrial application. In the stability evaluation, the used adsorbent was collected. Then, the stability test of the MHNSs was conducted by repeating the MO solution under the same conditions for five cycles without any regeneration treatment [33, 34].


## 3. Results and Discussion

### 3.1. XRD, Morphology, and the BET Measurements of the MHNSs

In this study, low-cost (NH4)2CO3 and KMnO4 were selected as the raw materials. The MHNSs were successfully fabricated by the selective etching method with the MnCO3 as the template. The typical SEM and TEM images of the MnCO3 template and MHNSs are shown in Figure 1. The MnCO3 templates show a well-defined microsphere structure with a diameter of about 2 μm, and the microspheres show a slight agglomeration between each other. Furthermore, some broken microspheres demonstrated that the spherical hollow structure of prepared MHNSs is derived from the selective etching of the MnCO3 template. The thickness of the spherical shell was about 300 nm. As shown in the high-resolution images, the original smooth nanosheets were replaced by the rough MnO2 nanoparticles. When the KMnO4 solution was added, the MnO2 nanoparticles formed by hydrolysis were adhered on the surface of the MnCO3 template. At the same time, the MnCO3 template was selectively etched by the HCl solution, then the MHNSs with a spherical hollow structure were obtained. Meanwhile, it can be observed that some dispersed MnO2 microsphere with the diameter at 200~300 nm coated on the surface of the hollow spheres. Furthermore, the hollow structure of prepared MHNSs was confirmed by the TEM techniques (Figure 1(g)–1(i)). The diameter of the hollow shells range from 1 μm to 2 μm, and the shell thickness was about 300 nm, which was consistent with the SEM analysis. Thus, a plausible formation mechanism was proposed in the scheme illustration figure (as shown in Figure 1).Figure 1
Scheme illustration and the typical SEM and TEM images of the samples.Figure2 showed the XRD patterns of prepared MnCO3 and MHNS samples. The diffraction peaks at 2θ=24.25°, 31.36°, and 51.68° can be assigned to the (012), (104), and (116) planes of MnCO3 (JCPDS card no. 44-1472), respectively. For the MHNS sample, the diffraction peaks at 2θ=37.12°, 45.01°, and 65.70° correspond to the (311), (400) and (440) crystal planes, respectively, which was consistent with the α-MnO2 (JCPDS card no. 44-0992). However, it can be observed that the MHNSs exhibit low and broad diffraction peaks. As mentioned in the FESEM discussions, the KMnO4 will hydrolyze to form MnO2 nanoparticles coated on the surface of the MnCO3 template. Meanwhile, the MnO2 maintains a low level of crystallinity due to the low-temperature and short-time reaction conditions. Thus, the characteristic peak intensity of the MHNSs was weak compared to that of prepared MnCO3 templates.Figure 2
XRD patterns of the MnCO3 template (A) and the MHNS samples (B).Figure 3
N2 adsorption-desorption isotherm (inset: the pore size distribution).In order to investigate the specific surface area, pore size distribution and pore volume of the MHNSs, the N2 adsorption-desorption isotherm was conducted and the result was as shown in Figure 3. The N2 adsorption-desorption isotherms exhibit type-IV isotherms. The BET surface area, total pore volume, and average pore size are determined to be 43.74 m2 g-1, 0.1133 cm3 g-1, and 1.04 nm, respectively. Thus, it can be concluded that the prepared MHNS sample is dominated by mesopores, promoting the adsorption ability for the MO molecule.


### 3.2. Effect of pH Value on MO Adsorption

Many studies have reported that the adsorption capacity of the adsorbent was pH-dependent. The reason can be that the pH value can influence the surface charge of the adsorbent and the degree of ionization/dissociation of the MO molecules. Thus, the adsorption on MO at different pH values are performed and shown in Figure4. With the increase of pH value from 2.0 to 11.0, the adsorption capacity of MHNSs on MO decreases from 99% to 2%. Obviously, MHNSs can only effectively adsorb MO molecules under acidic conditions, but it could hardly act at alkaline or neutral solution conditions. Since most of the free hydroxyl groups on the surface of MHNSs are derived from the hydrogen ions in the solution, the surface is positively charged. However, the sulfonate ions ionized by methyl orange molecules are negatively charged under acidic condition. Therefore, the negatively charged sulfonate ions will be adhered to the positively charged MHNS surface, then forming the ligand complex. At an acidic pH value, the solution is rich of hydrated hydrogen ions (H+ and H3O+), and the surface of MHNSs become more positively charged, which makes it more likely to adsorb MO in solution. On the contrary, the adsorption efficiency is relatively low under alkaline conditions. Moreover, there is a competitive adsorption relationship between the negatively charged MO ions and excessive hydroxide ions. With the increase of pH value, the concentration of hydroxide ions in the solution increases and the negative charge on the surface of MHNSs increases. So the Coulombic repulsion between the hydroxide ions and the anionic dye was enhanced, and the adsorption activity of MHNSs on the MO solution was decreased. Therefore, the significant influence of the adsorption amount of the MHNS surface was caused by the pH value of the solution, indicating that the electrostatic interaction may be the main mechanism for the adsorption process of anionic dyes. It can also be seen from Figure 5 that the adsorption process of MHNSs on MO is very fast, in which almost 90% of MO molecules were removed within 5 min.Figure 4
Effect of solution pH on the MO adsorption (experimental conditions: initial dye concentration: 50 mg L-1, adsorbent dosage: 20 mg/200 mL, temperature: 288 K, and contact time: 120 min).Figure 5
Time-dependent absorbance spectra of the adsorption process at selected intervals.Figure 6
Effect of contact time on MO adsorption at different initial concentrations.


### 3.3. Effect of Contact Time and Initial Concentration

Three initial concentrations (50, 100, and 150 mg L-1) of MO solutions were chosen to investigate the equilibrium time and the adsorption capacity of prepared MHNSs and the results was shown as in the Figure 6. With the increase of the initial concentration, the equilibrium absorption capacity (qe) of the MHNSs increased from 475.95 mg g-1 to 1420.08 mg g-1, indicating that the initial concentration had a significant impact on the adsorption capacity. The reason can be ascribed to the increasing driving force of concentration gradient. Moreover, the adsorption rate of MO was rapid at the initial 30 min and then slows (at the contact time of 30~120 min). During the initial stage, the MO molecules were rapidly absorbed by the MHNS adsorbent due to the sufficient adsorption sites on the surface. Then, the adsorption rate slowed down with the dwindling number of active sites and the increasing mutual repulsive forces between the MO molecules on the surface of MHNSs and the bulk solution. The reusability of the MHNS adsorbent was tested with the initial concentration of MO solution at 50 mg L-1, and the results are shown in Figure 7. It was found that the removal rate of MO was still maintained at 90% after four cycles. Therefore, it can be concluded that the MHNS adsorbent exhibited well adsorption stability on the MO solution.Figure 7
Efficiency of the MHNS composite after five cycles.


### 3.4. Evaluation of Adsorption Kinetics

The Lagergren-first-order equation, pseudo-second-order equation, and intraparticle diffusion model were used to investigate the adsorption kinetics of the MO molecule on MHNSs, and the correlation coefficient value (R2) was used to select the best-fit model [35, 36]. Figures 8–10 present the analysis results of MHNSs acted on the MO solution with different dynamic models. The pseudo-first-order kinetic model was specified as follows [37, 38]:
(2)lnqe−qt=lnqe−k1t,where qt (mg/g) is the adsorption capacity at time t (min), qe (mg/g) is the equilibrated adsorption capacity, and k1 (min−1) is the rate constant of the pseudo-first-order kinetic model.Figure 8
Regressions of kinetic plots at different initial concentrations: pseudo-first-order kinetic model.Figure 9
Regressions of kinetic plots at different initial concentrations: pseudo-second-order kinetic model.Figure 10
Regressions of kinetic plots at different initial concentrations: intraparticle mass transfer diffusion model.The value ofk1 can be calculated from the plots of lnqe‐qt verse t according to the equation. The linear formula of the pseudo-second-order kinetic model can be expressed as follows [39, 40]:
(3)tqt=1k2qe2+tqe,where k2 (g mg−1 min−1) is the rate constant of the pseudo-second-order kinetic model.The slope and intercept of the linear plots oft/qt against t was used to calculate the value of k2 and qe according to the equation.Then, the initial adsorption rateh (mg g-1 min-1) was determined as follows [41]:
(4)h=k2qe2.The actual rate-controlling step of the MO absorbed on the MHNSs was very crucial for researching the adsorption mechanism. However, the aforementioned two kinetic models could not predict it exactly, then the intraparticle mass transfer diffusion model proposed by Weber-Morris was employed. And the calculation equation was as follows [42, 43]:
(5)qt=kit1/2+C,where ki (mg g−1 min−1/2) is the intraparticle diffusion rate constant and C is the constant that reflects the thickness of the boundary layer.The graph ofqt against the square root of time can be used calculated the value of ki and the constant C.The three kinetic model parameters of MO adsorption onto MHNSs with different initial concentrations are shown in Table1. It could be observed that all the experimental data agreed well with the pseudo-second-order kinetic model (R2>0.9997). Therefore, the pseudo-second-order kinetic model was selected to analyze the adsorption process of MO molecule onto MHNSs in this study.Table 1
Parameters of pseudo-first- and pseudo-second-order adsorption kinetic models and the intraparticle diffusion model at various concentrations.
C0 (mg L-1)
qe,exp
Pseudo-first-order model
k1
qe
r2
50
475.95
0.04402
43.793
0.9869
100
942.99
0.04788
136.904
0.9683
150
1420.08
0.0598
560.220
0.9962
C0 (mg L-1)
Pseudo-second-order model
Intraparticle diffusion model
k2×10−3
qe
r2
h
ki
C
r2
50
2.7822
478.469
0.9999
636.943
8.339
421.176
0.9562
100
0.9187
952.381
0.9999
833.333
42.114
705.854
0.8299
150
0.2349
1458.99
0.9997
500.000
78.136
881.039
0.9830Generally, the adsorption process of MO on the MHNSs was controlled by the following three steps: external mass diffusion, intraparticle diffusion, and the adsorption on the MHNS surface. Because the adsorption process was very fast, this rate-limiting step often was not considered in this study. In order to determine the exact rate-controlling step, the intraparticle diffusion model was involved in the analysis process for MO adsorption on the MHNSs, and the results are shown in Figure10.Multivariate linear regression was used to analyze the MHNS adsorption MO data. According to the intraparticle diffusion model, the linear plot of adsorption capacity versus square rout of time (t1/2) indicated that there were two obvious regions in the graph. The first region was from 0 to 30 minutes, and the second one was 30-120 minutes. The linear portion in the first region has a larger slope, which represents the external mass transfer process. However, during the second stage, the adsorption capacity of MHNSs on MO still increased with t1/2, but compared with the first stage, the slope of the linear portion was much smaller. This can be ascribed to the second stage which was controlled by the intraparticle diffusion. Additionally, all the linear portions did not pass through the origin, indicating that the intraparticle diffusion was not the only rate-controlling step in the MHNS adsorption process. The ratio of the diffusion duration of the external mass to the intraparticle is about 1 : 3, which means that the whole adsorption process was controlled by the external mass transfer and the intraparticle diffusion, and the intraparticle diffusion was superior to the external mass transfer.


### 3.5. Adsorption Isotherm Study

In order to further study the adsorption interaction between the MO and the MHNSs, the equilibrium adsorption data were analyzed and fitted using the Langmuir and Freundlich isothermal models in this study [44, 45]. The Langmuir isotherm can be expressed as follows:
(6)CeQe=1QmKL+1QmCe,where Ce (mg/L) is the equilibrium concentration, Qe (mg/g) is the adsorption capacity of adsorbate per unit mass, Qm is the Langmuir constant, and KL is the adsorption rate.The essential characteristic of the Langmuir equation was expressed by the dimensionless separation factorRL:
(7)RL=11+KLC0,where KL (L/mg) is the Langmuir isotherm constant, C0 (mg/L) is the initial MO concentration, and RL is the type of isotherm.The linear form of the Freundlich equation is(8)lnQe=lnKF+1nlnCe,where KF (L/mg) is the adsorption capacity, and n is the definition of how favorable the adsorption process is.As shown in Table2, the RL value of the Langmuir model was between 0 and 1 and the n value of the Freundlich model was between 1 and 10. These results indicated that the adsorption process of MO molecules on the MHNSs was favorable. In addition, the correlation coefficient results of different models showed that the Freundlich model had a better fit of experimental data than that of the Langmuir model. When the temperature of the adsorption process increased from 273 K to 303 K, the maximum adsorption capacity of MO molecules on MHNSs decreased significantly from 1677.14 mg/g to 1589.40 mg/g. The maximum adsorption capacity of MHNSs and other adsorbents reported in literature on MO solution are listed in Table 3. It can be seen from the table that the saturated adsorption capacity of MHNSs was much higher than that of other adsorbents. Therefore, it can be considered that the prepared MHNSs have a potential application in the sewage treatment field.Table 2
Parameters of the adsorption isotherm for the removal of MO by the MHNSs at different temperatures.
T (K)
Langmuir
Freundlich
qm (mg g-1)
b
r2
RL
n
KF
r2
273
1718.712
0.2824
0.9992
0.01167~0.06614
1.6213
380.8478
0.9257
288
1688.416
0.2278
0.9989
0.01442~0.08071
1.1274
214.6052
0.9786
303
1633.290
0.2089
0.9981
0.01571~0.08737
1.1029
191.8282
0.9899Table 3
Comparison of the adsorption capacities of MO onto various adsorbents.
Adsorbates
Adsorbents
qmax (mg g-1)
Reference
MO
Amine modified PIM-1 ultrafine fibers
312.50 mg g−1
[46]
3-dimensional (3D) flower-like Ni/Al@PAB LDH
412.80 mg g-1
[47]
Date palm ash (DPA) and MgAl-LDH composites
242.98 mg g-1
[48]
Magnetic lignin-based carbon nanoparticles (MLBCN)
113.00 mg g-1
[49]
Chitosan-lysozyme biocomposite (CLC)
435.00 mg g-1
[50]
MnO2-graphene-carbon nanotube (MnO2-G-CNT)
476.19 mg g-1
[51]
CaFe2O4 and ZrO2 magnetic nanocomposite (CaF-ZO-MNC)
370.37 mg g-1
[52]
An amorphous solid contains large cationic aMOC-1
359.00 mg g-1
[53]
β-CD and PEI bifunctionalized magnetic (Fe3O4-PEI/β-CD)
192.20 mg g-1
[54]
The MnO2 hollow nanospheres (MHNSs)
1677.13 mg g-1
This work


### 3.6. Thermomechanical Analysis

In general, depth information such as the internal energy transformation can be analyzed by the thermodynamic parameters in the adsorption process. The standard Gibbs free energy changeΔG°, the standard molar enthalpy change ΔH° and the standard molar entropy change ΔS° were determined according to the adsorption thermodynamic equilibrium constants at different temperatures. For the adsorption process, the adsorption equilibrium constant K0 can be defined by the following formula [55, 56]:
(9)K=αsαe=γsγe⋅qeCe,where αs is the activity (effective concentration) of adsorbed methyl orange and αe is the activity of methyl orange in the equilibrium solution. γs is the activity coefficient of adsorption methyl orange. γe is the activity coefficient of methyl orange in equilibrium solution. With the progress of adsorption, the concentration of methyl orange in the solution keeps decreasing and tends to zero when the adsorption equilibrium is reached. Therefore, the equilibrium constant K0 value could be obtained by plotting lnqe/Ce versus qe and extrapolating qe equal to zero.For any interaction, the adsorption standard Gibbs free energy changeΔG° can be expressed by
(10)ΔG°=−RTlnK0,ΔG°=ΔH°−TΔS°.From the above two formulas,K0 can be concluded as follows:
(11)lnK0=−ΔH°RT+ΔS°R,where R is the standard gas constant (8.314 J mol-1 K-1) and T is the Kelvin temperature (K); ΔH° and ΔS° can be calculated from the slope and intercept of plot lnK0 versus 1/T. The thermodynamic data is listed in Table 4. It can be seen that the ΔG° values at the three temperatures were all negative, indicating that the adsorption of MO on the surface of MHNSs was a spontaneous process and thermodynamically feasible. In general, the ΔG° of the physical adsorption process was less than that of the chemical adsorption process. Normally, the ΔG° value of the physical adsorption process was between -20 and 0 kJ mol-1, while that of the chemisorption process was between -400 and -80 kJ mol-1 [57]. And the adsorption standard value ΔH° of the physical process was 2~21 kJ mol-1 [58]. Therefore, the value of ΔH° (-16.825) also suggested that the adsorption of MO onto the surface of MHNSs was mainly a physical process, which was also consistent with the previous results. Therefore, the adsorption process of MO on MHNSs was mainly a physical process. Moreover, the negative value of ΔH° revealed that the adsorption process is an exothermic process, which was well consistent with the aforementioned results that the adsorption capacity of MO decreased with the increase of temperature. Meanwhile, the negative ΔS° of the adsorption demonstrated that the MO molecule in bulk solution was in a more chaotic distribution compared to the ordered state (adsorbed on the surface of MHNSs).Table 4
Values of thermodynamic parameters for the adsorption of the MO dye on the MHNSs.
T (K)
lnK0
ΔG° (kJ mol-1)
ΔH° (kJ mol-1)
ΔS° (J mol-1 K-1)
273
7.3756
-16.740
-16.825
-1.160
288
6.7616
-16.189
303
6.6528
-16.759


### 3.7. Adsorption Mechanism Analysis

In order to further study the mechanism of the MO adsorption process on adsorbents, the FT-IR spectra of the fresh and used MHNS samples are shown in Figure11. The weak peaks around at about 530 cm-1 in the low frequency region could be attributed to the Mn-O and Mn-O-Mn shrink vibration [59, 60]. Importantly, the FT-IR spectra of the MHNS-MO sample were significantly different from its original spectrum. In the case of the MHNS-MO sample, the new peaks located at 1039, 1193, and 1384 cm-1 correspond to the contraction vibrations of the -S=O, CH3-, and CN bonds, respectively. And the peaks located at 818, 849 cm-1 can be indexed to the fingerprint region indicating that MO molecules were concentrated on the surface of MHNSs [61, 62].Figure 11
FT-IR spectra of pure MO and the MHNS composites before and after adsorption of MO.Generally, adsorption reaction may cause some changes of physical properties changes in the adsorbent, and the resulting change may help to investigate the adsorption reaction. As shown in Figure12, the surface chemical composition and valence state of MHNSs before and after adsorption of MO molecules were characterized by XPS. It can be seen that the peaks of Mn 2p3/2 and Mn 2p1/2 were located at 642.1 eV and 653.5 eV, respectively, indicating that the Mn of the prepared MHNSs were in the form of Mn4+. The energy separation between the two peaks was 11.4 eV, which was well in agreement with the previous literatures [63, 64]. Furthermore, the position and intensity of the Mn 2p peaks did not change after the adsorption, which confirmed that the adsorption process of MO did not cause a significant change in the oxidation state of manganese. Additionally, it was found that there was a weak N 1s peak located at 399.8 eV for the adsorbed MHNS-MO spectrum, which indicated that a large amount of MO was adsorbed onto the MHNSs. These results indicated that the adsorption of MO molecules did not change the chemical composition and valence states of MHNSs. This adsorption process was mainly a physical process, which was consistent with the conclusions in thermomechanical analysis.Figure 12
XPS spectra of the MHNS composites before and after adsorption of MO.


## 3.1. XRD, Morphology, and the BET Measurements of the MHNSs

In this study, low-cost (NH4)2CO3 and KMnO4 were selected as the raw materials. The MHNSs were successfully fabricated by the selective etching method with the MnCO3 as the template. The typical SEM and TEM images of the MnCO3 template and MHNSs are shown in Figure 1. The MnCO3 templates show a well-defined microsphere structure with a diameter of about 2 μm, and the microspheres show a slight agglomeration between each other. Furthermore, some broken microspheres demonstrated that the spherical hollow structure of prepared MHNSs is derived from the selective etching of the MnCO3 template. The thickness of the spherical shell was about 300 nm. As shown in the high-resolution images, the original smooth nanosheets were replaced by the rough MnO2 nanoparticles. When the KMnO4 solution was added, the MnO2 nanoparticles formed by hydrolysis were adhered on the surface of the MnCO3 template. At the same time, the MnCO3 template was selectively etched by the HCl solution, then the MHNSs with a spherical hollow structure were obtained. Meanwhile, it can be observed that some dispersed MnO2 microsphere with the diameter at 200~300 nm coated on the surface of the hollow spheres. Furthermore, the hollow structure of prepared MHNSs was confirmed by the TEM techniques (Figure 1(g)–1(i)). The diameter of the hollow shells range from 1 μm to 2 μm, and the shell thickness was about 300 nm, which was consistent with the SEM analysis. Thus, a plausible formation mechanism was proposed in the scheme illustration figure (as shown in Figure 1).Figure 1
Scheme illustration and the typical SEM and TEM images of the samples.Figure2 showed the XRD patterns of prepared MnCO3 and MHNS samples. The diffraction peaks at 2θ=24.25°, 31.36°, and 51.68° can be assigned to the (012), (104), and (116) planes of MnCO3 (JCPDS card no. 44-1472), respectively. For the MHNS sample, the diffraction peaks at 2θ=37.12°, 45.01°, and 65.70° correspond to the (311), (400) and (440) crystal planes, respectively, which was consistent with the α-MnO2 (JCPDS card no. 44-0992). However, it can be observed that the MHNSs exhibit low and broad diffraction peaks. As mentioned in the FESEM discussions, the KMnO4 will hydrolyze to form MnO2 nanoparticles coated on the surface of the MnCO3 template. Meanwhile, the MnO2 maintains a low level of crystallinity due to the low-temperature and short-time reaction conditions. Thus, the characteristic peak intensity of the MHNSs was weak compared to that of prepared MnCO3 templates.Figure 2
XRD patterns of the MnCO3 template (A) and the MHNS samples (B).Figure 3
N2 adsorption-desorption isotherm (inset: the pore size distribution).In order to investigate the specific surface area, pore size distribution and pore volume of the MHNSs, the N2 adsorption-desorption isotherm was conducted and the result was as shown in Figure 3. The N2 adsorption-desorption isotherms exhibit type-IV isotherms. The BET surface area, total pore volume, and average pore size are determined to be 43.74 m2 g-1, 0.1133 cm3 g-1, and 1.04 nm, respectively. Thus, it can be concluded that the prepared MHNS sample is dominated by mesopores, promoting the adsorption ability for the MO molecule.


## 3.2. Effect of pH Value on MO Adsorption

Many studies have reported that the adsorption capacity of the adsorbent was pH-dependent. The reason can be that the pH value can influence the surface charge of the adsorbent and the degree of ionization/dissociation of the MO molecules. Thus, the adsorption on MO at different pH values are performed and shown in Figure4. With the increase of pH value from 2.0 to 11.0, the adsorption capacity of MHNSs on MO decreases from 99% to 2%. Obviously, MHNSs can only effectively adsorb MO molecules under acidic conditions, but it could hardly act at alkaline or neutral solution conditions. Since most of the free hydroxyl groups on the surface of MHNSs are derived from the hydrogen ions in the solution, the surface is positively charged. However, the sulfonate ions ionized by methyl orange molecules are negatively charged under acidic condition. Therefore, the negatively charged sulfonate ions will be adhered to the positively charged MHNS surface, then forming the ligand complex. At an acidic pH value, the solution is rich of hydrated hydrogen ions (H+ and H3O+), and the surface of MHNSs become more positively charged, which makes it more likely to adsorb MO in solution. On the contrary, the adsorption efficiency is relatively low under alkaline conditions. Moreover, there is a competitive adsorption relationship between the negatively charged MO ions and excessive hydroxide ions. With the increase of pH value, the concentration of hydroxide ions in the solution increases and the negative charge on the surface of MHNSs increases. So the Coulombic repulsion between the hydroxide ions and the anionic dye was enhanced, and the adsorption activity of MHNSs on the MO solution was decreased. Therefore, the significant influence of the adsorption amount of the MHNS surface was caused by the pH value of the solution, indicating that the electrostatic interaction may be the main mechanism for the adsorption process of anionic dyes. It can also be seen from Figure 5 that the adsorption process of MHNSs on MO is very fast, in which almost 90% of MO molecules were removed within 5 min.Figure 4
Effect of solution pH on the MO adsorption (experimental conditions: initial dye concentration: 50 mg L-1, adsorbent dosage: 20 mg/200 mL, temperature: 288 K, and contact time: 120 min).Figure 5
Time-dependent absorbance spectra of the adsorption process at selected intervals.Figure 6
Effect of contact time on MO adsorption at different initial concentrations.


## 3.3. Effect of Contact Time and Initial Concentration

Three initial concentrations (50, 100, and 150 mg L-1) of MO solutions were chosen to investigate the equilibrium time and the adsorption capacity of prepared MHNSs and the results was shown as in the Figure 6. With the increase of the initial concentration, the equilibrium absorption capacity (qe) of the MHNSs increased from 475.95 mg g-1 to 1420.08 mg g-1, indicating that the initial concentration had a significant impact on the adsorption capacity. The reason can be ascribed to the increasing driving force of concentration gradient. Moreover, the adsorption rate of MO was rapid at the initial 30 min and then slows (at the contact time of 30~120 min). During the initial stage, the MO molecules were rapidly absorbed by the MHNS adsorbent due to the sufficient adsorption sites on the surface. Then, the adsorption rate slowed down with the dwindling number of active sites and the increasing mutual repulsive forces between the MO molecules on the surface of MHNSs and the bulk solution. The reusability of the MHNS adsorbent was tested with the initial concentration of MO solution at 50 mg L-1, and the results are shown in Figure 7. It was found that the removal rate of MO was still maintained at 90% after four cycles. Therefore, it can be concluded that the MHNS adsorbent exhibited well adsorption stability on the MO solution.Figure 7
Efficiency of the MHNS composite after five cycles.


## 3.4. Evaluation of Adsorption Kinetics

The Lagergren-first-order equation, pseudo-second-order equation, and intraparticle diffusion model were used to investigate the adsorption kinetics of the MO molecule on MHNSs, and the correlation coefficient value (R2) was used to select the best-fit model [35, 36]. Figures 8–10 present the analysis results of MHNSs acted on the MO solution with different dynamic models. The pseudo-first-order kinetic model was specified as follows [37, 38]:
(2)lnqe−qt=lnqe−k1t,where qt (mg/g) is the adsorption capacity at time t (min), qe (mg/g) is the equilibrated adsorption capacity, and k1 (min−1) is the rate constant of the pseudo-first-order kinetic model.Figure 8
Regressions of kinetic plots at different initial concentrations: pseudo-first-order kinetic model.Figure 9
Regressions of kinetic plots at different initial concentrations: pseudo-second-order kinetic model.Figure 10
Regressions of kinetic plots at different initial concentrations: intraparticle mass transfer diffusion model.The value ofk1 can be calculated from the plots of lnqe‐qt verse t according to the equation. The linear formula of the pseudo-second-order kinetic model can be expressed as follows [39, 40]:
(3)tqt=1k2qe2+tqe,where k2 (g mg−1 min−1) is the rate constant of the pseudo-second-order kinetic model.The slope and intercept of the linear plots oft/qt against t was used to calculate the value of k2 and qe according to the equation.Then, the initial adsorption rateh (mg g-1 min-1) was determined as follows [41]:
(4)h=k2qe2.The actual rate-controlling step of the MO absorbed on the MHNSs was very crucial for researching the adsorption mechanism. However, the aforementioned two kinetic models could not predict it exactly, then the intraparticle mass transfer diffusion model proposed by Weber-Morris was employed. And the calculation equation was as follows [42, 43]:
(5)qt=kit1/2+C,where ki (mg g−1 min−1/2) is the intraparticle diffusion rate constant and C is the constant that reflects the thickness of the boundary layer.The graph ofqt against the square root of time can be used calculated the value of ki and the constant C.The three kinetic model parameters of MO adsorption onto MHNSs with different initial concentrations are shown in Table1. It could be observed that all the experimental data agreed well with the pseudo-second-order kinetic model (R2>0.9997). Therefore, the pseudo-second-order kinetic model was selected to analyze the adsorption process of MO molecule onto MHNSs in this study.Table 1
Parameters of pseudo-first- and pseudo-second-order adsorption kinetic models and the intraparticle diffusion model at various concentrations.
C0 (mg L-1)
qe,exp
Pseudo-first-order model
k1
qe
r2
50
475.95
0.04402
43.793
0.9869
100
942.99
0.04788
136.904
0.9683
150
1420.08
0.0598
560.220
0.9962
C0 (mg L-1)
Pseudo-second-order model
Intraparticle diffusion model
k2×10−3
qe
r2
h
ki
C
r2
50
2.7822
478.469
0.9999
636.943
8.339
421.176
0.9562
100
0.9187
952.381
0.9999
833.333
42.114
705.854
0.8299
150
0.2349
1458.99
0.9997
500.000
78.136
881.039
0.9830Generally, the adsorption process of MO on the MHNSs was controlled by the following three steps: external mass diffusion, intraparticle diffusion, and the adsorption on the MHNS surface. Because the adsorption process was very fast, this rate-limiting step often was not considered in this study. In order to determine the exact rate-controlling step, the intraparticle diffusion model was involved in the analysis process for MO adsorption on the MHNSs, and the results are shown in Figure10.Multivariate linear regression was used to analyze the MHNS adsorption MO data. According to the intraparticle diffusion model, the linear plot of adsorption capacity versus square rout of time (t1/2) indicated that there were two obvious regions in the graph. The first region was from 0 to 30 minutes, and the second one was 30-120 minutes. The linear portion in the first region has a larger slope, which represents the external mass transfer process. However, during the second stage, the adsorption capacity of MHNSs on MO still increased with t1/2, but compared with the first stage, the slope of the linear portion was much smaller. This can be ascribed to the second stage which was controlled by the intraparticle diffusion. Additionally, all the linear portions did not pass through the origin, indicating that the intraparticle diffusion was not the only rate-controlling step in the MHNS adsorption process. The ratio of the diffusion duration of the external mass to the intraparticle is about 1 : 3, which means that the whole adsorption process was controlled by the external mass transfer and the intraparticle diffusion, and the intraparticle diffusion was superior to the external mass transfer.


## 3.5. Adsorption Isotherm Study

In order to further study the adsorption interaction between the MO and the MHNSs, the equilibrium adsorption data were analyzed and fitted using the Langmuir and Freundlich isothermal models in this study [44, 45]. The Langmuir isotherm can be expressed as follows:
(6)CeQe=1QmKL+1QmCe,where Ce (mg/L) is the equilibrium concentration, Qe (mg/g) is the adsorption capacity of adsorbate per unit mass, Qm is the Langmuir constant, and KL is the adsorption rate.The essential characteristic of the Langmuir equation was expressed by the dimensionless separation factorRL:
(7)RL=11+KLC0,where KL (L/mg) is the Langmuir isotherm constant, C0 (mg/L) is the initial MO concentration, and RL is the type of isotherm.The linear form of the Freundlich equation is(8)lnQe=lnKF+1nlnCe,where KF (L/mg) is the adsorption capacity, and n is the definition of how favorable the adsorption process is.As shown in Table2, the RL value of the Langmuir model was between 0 and 1 and the n value of the Freundlich model was between 1 and 10. These results indicated that the adsorption process of MO molecules on the MHNSs was favorable. In addition, the correlation coefficient results of different models showed that the Freundlich model had a better fit of experimental data than that of the Langmuir model. When the temperature of the adsorption process increased from 273 K to 303 K, the maximum adsorption capacity of MO molecules on MHNSs decreased significantly from 1677.14 mg/g to 1589.40 mg/g. The maximum adsorption capacity of MHNSs and other adsorbents reported in literature on MO solution are listed in Table 3. It can be seen from the table that the saturated adsorption capacity of MHNSs was much higher than that of other adsorbents. Therefore, it can be considered that the prepared MHNSs have a potential application in the sewage treatment field.Table 2
Parameters of the adsorption isotherm for the removal of MO by the MHNSs at different temperatures.
T (K)
Langmuir
Freundlich
qm (mg g-1)
b
r2
RL
n
KF
r2
273
1718.712
0.2824
0.9992
0.01167~0.06614
1.6213
380.8478
0.9257
288
1688.416
0.2278
0.9989
0.01442~0.08071
1.1274
214.6052
0.9786
303
1633.290
0.2089
0.9981
0.01571~0.08737
1.1029
191.8282
0.9899Table 3
Comparison of the adsorption capacities of MO onto various adsorbents.
Adsorbates
Adsorbents
qmax (mg g-1)
Reference
MO
Amine modified PIM-1 ultrafine fibers
312.50 mg g−1
[46]
3-dimensional (3D) flower-like Ni/Al@PAB LDH
412.80 mg g-1
[47]
Date palm ash (DPA) and MgAl-LDH composites
242.98 mg g-1
[48]
Magnetic lignin-based carbon nanoparticles (MLBCN)
113.00 mg g-1
[49]
Chitosan-lysozyme biocomposite (CLC)
435.00 mg g-1
[50]
MnO2-graphene-carbon nanotube (MnO2-G-CNT)
476.19 mg g-1
[51]
CaFe2O4 and ZrO2 magnetic nanocomposite (CaF-ZO-MNC)
370.37 mg g-1
[52]
An amorphous solid contains large cationic aMOC-1
359.00 mg g-1
[53]
β-CD and PEI bifunctionalized magnetic (Fe3O4-PEI/β-CD)
192.20 mg g-1
[54]
The MnO2 hollow nanospheres (MHNSs)
1677.13 mg g-1
This work


## 3.6. Thermomechanical Analysis

In general, depth information such as the internal energy transformation can be analyzed by the thermodynamic parameters in the adsorption process. The standard Gibbs free energy changeΔG°, the standard molar enthalpy change ΔH° and the standard molar entropy change ΔS° were determined according to the adsorption thermodynamic equilibrium constants at different temperatures. For the adsorption process, the adsorption equilibrium constant K0 can be defined by the following formula [55, 56]:
(9)K=αsαe=γsγe⋅qeCe,where αs is the activity (effective concentration) of adsorbed methyl orange and αe is the activity of methyl orange in the equilibrium solution. γs is the activity coefficient of adsorption methyl orange. γe is the activity coefficient of methyl orange in equilibrium solution. With the progress of adsorption, the concentration of methyl orange in the solution keeps decreasing and tends to zero when the adsorption equilibrium is reached. Therefore, the equilibrium constant K0 value could be obtained by plotting lnqe/Ce versus qe and extrapolating qe equal to zero.For any interaction, the adsorption standard Gibbs free energy changeΔG° can be expressed by
(10)ΔG°=−RTlnK0,ΔG°=ΔH°−TΔS°.From the above two formulas,K0 can be concluded as follows:
(11)lnK0=−ΔH°RT+ΔS°R,where R is the standard gas constant (8.314 J mol-1 K-1) and T is the Kelvin temperature (K); ΔH° and ΔS° can be calculated from the slope and intercept of plot lnK0 versus 1/T. The thermodynamic data is listed in Table 4. It can be seen that the ΔG° values at the three temperatures were all negative, indicating that the adsorption of MO on the surface of MHNSs was a spontaneous process and thermodynamically feasible. In general, the ΔG° of the physical adsorption process was less than that of the chemical adsorption process. Normally, the ΔG° value of the physical adsorption process was between -20 and 0 kJ mol-1, while that of the chemisorption process was between -400 and -80 kJ mol-1 [57]. And the adsorption standard value ΔH° of the physical process was 2~21 kJ mol-1 [58]. Therefore, the value of ΔH° (-16.825) also suggested that the adsorption of MO onto the surface of MHNSs was mainly a physical process, which was also consistent with the previous results. Therefore, the adsorption process of MO on MHNSs was mainly a physical process. Moreover, the negative value of ΔH° revealed that the adsorption process is an exothermic process, which was well consistent with the aforementioned results that the adsorption capacity of MO decreased with the increase of temperature. Meanwhile, the negative ΔS° of the adsorption demonstrated that the MO molecule in bulk solution was in a more chaotic distribution compared to the ordered state (adsorbed on the surface of MHNSs).Table 4
Values of thermodynamic parameters for the adsorption of the MO dye on the MHNSs.
T (K)
lnK0
ΔG° (kJ mol-1)
ΔH° (kJ mol-1)
ΔS° (J mol-1 K-1)
273
7.3756
-16.740
-16.825
-1.160
288
6.7616
-16.189
303
6.6528
-16.759


## 3.7. Adsorption Mechanism Analysis

In order to further study the mechanism of the MO adsorption process on adsorbents, the FT-IR spectra of the fresh and used MHNS samples are shown in Figure11. The weak peaks around at about 530 cm-1 in the low frequency region could be attributed to the Mn-O and Mn-O-Mn shrink vibration [59, 60]. Importantly, the FT-IR spectra of the MHNS-MO sample were significantly different from its original spectrum. In the case of the MHNS-MO sample, the new peaks located at 1039, 1193, and 1384 cm-1 correspond to the contraction vibrations of the -S=O, CH3-, and CN bonds, respectively. And the peaks located at 818, 849 cm-1 can be indexed to the fingerprint region indicating that MO molecules were concentrated on the surface of MHNSs [61, 62].Figure 11
FT-IR spectra of pure MO and the MHNS composites before and after adsorption of MO.Generally, adsorption reaction may cause some changes of physical properties changes in the adsorbent, and the resulting change may help to investigate the adsorption reaction. As shown in Figure12, the surface chemical composition and valence state of MHNSs before and after adsorption of MO molecules were characterized by XPS. It can be seen that the peaks of Mn 2p3/2 and Mn 2p1/2 were located at 642.1 eV and 653.5 eV, respectively, indicating that the Mn of the prepared MHNSs were in the form of Mn4+. The energy separation between the two peaks was 11.4 eV, which was well in agreement with the previous literatures [63, 64]. Furthermore, the position and intensity of the Mn 2p peaks did not change after the adsorption, which confirmed that the adsorption process of MO did not cause a significant change in the oxidation state of manganese. Additionally, it was found that there was a weak N 1s peak located at 399.8 eV for the adsorbed MHNS-MO spectrum, which indicated that a large amount of MO was adsorbed onto the MHNSs. These results indicated that the adsorption of MO molecules did not change the chemical composition and valence states of MHNSs. This adsorption process was mainly a physical process, which was consistent with the conclusions in thermomechanical analysis.Figure 12
XPS spectra of the MHNS composites before and after adsorption of MO.


## 4. Conclusions

The MnO2 hollow nanospheres (MHNSs) were prepared through a selective etching method using MnCO3 as sacrificial templates. The crystal structure, morphology, and BET surface area were investigated. The adsorption capacity of MO onto the MHNSs was systematically investigated by adjusting the pH value, the initial MO concentration, and the contact time. The prepared MHNSs showed an excellent adsorption capacity toward MO compared to other adsorbents. The adsorption kinetics followed the pseudo-second-order model. Moreover, the overall process was apparently controlled by the external mass transfer and intraparticle diffusion based on the adsorption kinetics data. The thermodynamic analysis showed that the adsorption of MO on MHNSs was a spontaneous, exothermic, and physical process. Similarly, the XPS results also confirmed that the adsorption was mainly a physical process. The superior adsorbing performance of the MHNSs could be attributed to their unique hollow structure and large BET surface area, which provided much active adsorption sites. It is expected that the prepared MHNSs may be a promising adsorbent for water treatment due to its superior adsorbing performance and good stability.



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