Environmental Science and Pollution Research

, Volume 26, Issue 18, pp 18825–18833 | Cite as

Facile synthesis of hollow mesoporous MgO spheres via spray-drying with improved adsorption capacity for Pb(II) and Cd(II)

  • Mengjie Kuang
  • Yisheng Shang
  • Gaoling Yang
  • Baixiong LiuEmail author
  • Bin YangEmail author
Research Article


Spherical-like MgO nanostructures have been synthesized efficiently via spray-drying combined with calcination using magnesium acetate as magnesium source. The products were characterized by means of X-ray powder diffraction (XRD), transmission electron microscopy (TEM), scanning electron microscopy (SEM), and the specific surface areas were calculated using the Brunauer–Emmett–Teller (BET) method. The obtained spherical-like MgO nanostructures exhibit uniform pore sizes (7.7 nm) and high specific surface areas (180 m2 g−1). The adsorption kinetics and isotherm data agree well with pseudo-second-order model and Langmuir model, indicating the monolayer chemisorption of heavy metal ions. The spherical-like MgO nanostructures exhibited high adsorption performance for Pb(II) and Cd(II) ions, and the maximum adsorption capacities were up to 5214 mg g−1 and 4187 mg g−1, respectively. These values are much higher than those reported MgO-based adsorbents. Moreover, in less than 10 min, Pb(II) and Cd(II) ions in solution can be almost removed, which means that the spherical-like MgO possesses a high adsorption rate. XRD and FTIR analysis revealed the adsorption mechanism of Pb(II) and Cd(II) ions on MgO, which was mainly due to hydroxyl functional groups and ion exchange between Mg and heavy metal ions on the surface of MgO. These favorable performances recommend that the synthesized spherical-like MgO nanostructures would be a potential adsorbent for rapid removal of heavy metal ions from wastewater.


Spherical-like MgO Nanostructures Magnesium oxide Spray-drying Adsorption Heavy metal ions 


Water pollution by heavy metal ions, which have toxic effects on human health, as skin disease, respiratory system disorders, heart disease, and kidney or liver damage, is becoming a serious problem (Qu et al. 2013). Consequently, methods for effective removal of such toxic pollutants from water are urgently required and will have great significance in wastewater purification (Chowdhury et al. 2016; Jia et al. 2013; Wang et al. 2010). Up to now, various treatments have been developed to remove heavy metal ions such as electrochemical method, ion exchange, chemical coagulation, oxidation, adsorption, membrane separation, and so on (Brillas and Martínez-Huitle 2009; Giustetto et al. 2012; Iannazzo et al. 2017; Li and Zhou 2018; Mansouri et al. 2011; Marto et al. 2009). Among these methods, adsorption has been generally considered as one of the most promising techniques for elimination of heavy metal ions from the standpoints of simple and stable operation process, high efficiency, and low cost (Li et al. 2012). Many adsorbents, including zeolite, carbon, and activated alumina, etc., are widely used to adsorb heavy metal ions in aqueous solution. However, many factors have restricted their applications, such as the limited adsorption abilities, complex synthetic method and high operating cost. Therefore, it is necessary to explore adsorbents that efficiently remove heavy metal ions. In recent years, due to unique properties of high surface area, environmentally benign nature, and low production cost, some metal oxides with nanostructures as absorbents have attracted much attention. Nanostructure strontium oxide, copper oxide, and iron oxide have been used to adsorb metal ions from wastewater. Nevertheless, the adsorption capacities and removal efficiencies of these oxides are still great challenges. Hence, it is highly desirable to develop effective, high adsorption performance and reliable nanostructure metal oxides which can resolve the problems of wastewater treatment.

Nanostructured MgO as a typical metal oxide has been already used as adsorbent for removing heavy metal ions from wastewater due to its unique physical and chemical properties, such as high surface areas, a lot of active sites, and the favorable electrostatic attraction mechanism (Alvarado et al. 2000; Kumar and Kumar 2008; Khairallah and Glisenti 2007). Furthermore, nanostructured MgO is also nontoxic, thermally stable, environment friendly, economical, and easy to synthesize. These properties recommend that nanostructured MgO is an economical and effective adsorbent that can rapidly remove heavy metal ions from water. Thus, nanostructured MgO is considered as one of the most ideal adsorbents for water treatment.

At present, many methods for synthesizing MgO have been reported, such as sol-gel, aqueous wet chemical, hydrothermal, microwave-assisted, flame spray pyrolysis, combustion aerosol synthesis, chemical gas phase deposition, and surfactant method etc. (Kuang et al. 2018; Kaithik et al. 2019; Possato et al. 2018; Tang et al. 2018). Among these methods, hydrothermal method is relatively widely used to synthesize MgO nanostructures with different shapes (Jia et al. 2014; Li et al. 2012; She et al. 2011; Yu et al. 2011). Chowdhury et al. (2016) synthesized mesoporous MgO nanostructures by hydrothermal method at 180 °C for 5 h and calcined at 600 °C by using different magnesium salts without any templating agent. The as-synthesized MgO exhibited maximum adsorption capacities up to 2550 mg g−1 for Pb(II) and 2491 mg g−1 for Cd(II). Microcubes of MgO-TiO2 have been synthesized by hydrothermal method at 180 °C and calcined at 650 °C by using magnesium acetate and titanium (IV) oxysulfate (TIOS) as MgO precursor and the source of TiO2. The prepared MgO–TiO2 microcubes showed high adsorption property for Pb (II) ions, with capacity of 2900 mg g−1 (Xiong et al. 2015). However, these methods have their own demerits in the preparation of nanostructured MgO, such as complex synthesis process and low efficiency.

Spray-drying is a method that prepares powders by dispersing the solution into fog droplets with high-pressure gas and drying these droplets through hot air. Therefore, spray-drying method can achieve a rapid transfer between heat and mass in a very short time. These properties indicate that spray-drying is a method with low cost and high efficiency, which would be suitable for industrialization production. In the present study, we demonstrated a method to synthesize spherical-like MgO by spray-drying combined with calcination with improved adsorption activities and capacities.



Magnesium acetate (Mg(CH3COO)2·4H2O), ethanol, lead nitrate (Pb(NO3)2·4H2O), and cadmium nitrate(Cd(NO3)2·4H2O) are of analytical grade, and they were purchased from Beijing Chemicals Reagent Company (Beijing, China). All chemicals were used without any further purification.

Synthesis of spherical-like MgO nanostructures

Forty grams Mg(CH3COO)2·4H2O was dissolved in deionized water to prepare 40 g L−1 magnesium acetate aqueous solution. Subsequently, pure precursor powders were obtained by spray-drying at 230 °C and the feeding rate of 5 mL·min−1. Finally, the powder products were obtained by calcination at 350 °C for 1.5 h at a rate of 0.1 °C/min in a rotary tube furnace.


The thermogravimetry (TG) and differential thermal analysis (DTA) (TG-DTA 6300, Perkin Elmer) were used to analyze the thermal behavior of the samples. The measurements were conducted at a heating rate of 10 °C min−1 from 20 °C to 1000 °C in air atmosphere. X-ray diffraction (XRD) patterns of the specimen were recorded on XRD Rigaku D/max-2500 using Cu Kα (λ = 0.15 nm) radiation with a scanning speed 3 min−1. The morphologies were studied by using field emission scanning electron microscope (FESEM) (HITACHI S4800) with work voltage of 15 kV. The transmission electron microscopy (TEM) (Tecnai G2 30ST) was utilized to observe the high-resolution morphologies, with an accelerating voltage of 300 kV. N2 adsorption and desorption isotherms were measured at 77 K by using Micromeritics ASAP 2020. The Brunauer–Emmett–Teller (BET) and Barrett–Joyner–Halenda (BJH) methods were used to determine the specific surface areas and the pore size distributions. The concentrations of Pb(II) and Cd(II) ions remaining in solution after adsorption were determined by using inductively coupled plasma-atomic emission spectroscopy (ICP-AES, Jarrell-Ash model ICAP 9000).

Heavy metal ion adsorption test

In the experiments, Pb(NO3)2·4H2O and Cd(NO3)2·4H2O were used as the sources of Pb(II) and Cd(II), respectively. All pH values in the solutions with different concentrations of Pb(II) and Cd(II) ions were 5. The as-prepared MgO spheres were used as adsorbents. For the adsorption isotherm study, 30 mg of MgO was added into 30 mL Pb(II) and Cd(II) solutions with different concentrations under stirring at room temperature. After a certain time, the solid and liquid were separated immediately, and then used the inductively coupled plasma-atomic emission spectroscopy (ICP) to measure the concentration of Pb(II) and Cd(II) ions after adsorption. For the adsorption kinetic study, 240 mg of MgO powders was added into 240 mL solutions with 100 mg/L Pb(II) and Cd(II), respectively. The solutions were stirring with a specified time, then separated the solid and liquid immediately, and ICP was used to measure the Pb(II) or Cd(II) concentrations remaining in the solutions.

Results and discussion

The thermal behavior of the obtained Mg(CH3COO)2·4H2O was analyzed by TG-DTA technique. According to our previous experimental results, the calcination temperature should be set up to 328.8 °C for ensuring Mg(CH3COO)2 was completely decomposed (Kuang et al. 2018).

The powders were easily prepared via spray-drying and thermal treatment. The phase structures of final products and precursors were characterized by XRD. The X-ray diffraction curve of the precursors displays a broad and weak reflection at 2θ of 25°, which is typical characteristic of nanocrystalline components (as shown in Fig. 1a). This result indicated that the precursors were amorphous. The XRD pattern of the samples obtained by calcination of the precursors at 350 °C for 1.5 h is shown in Fig. 1b. All diffraction peaks at 2θ of 36.9° (111), 42.9° (200), 62.3° (220), 74.6° (311), 78.6° (222) can be identified as cubic MgO according to JCPDS no. 45–0946. Moreover, there were no impurity peaks, indicating that the products were pure MgO nanoparticles.
Fig. 1

XRD patterns of a as-prepared precursors and b as-prepared samples after calcined

We studied the effect of heating rate on the morphology of MgO products during calcination. As shown in Fig. 2a, the final products show flower-like morphologies formed by twisted nanosheets. Figure 2b shows the detailed morphology, it can be observed that the entire structure consists of nanosheets with smooth surfaces. While in Fig. 2c, the final products show spherical-like morphologies, the surfaces and edges of the spherical particles are smooth. The high-magnification SEM image in Fig. 2d can be clearly observed that the prepared samples have hollow structures through the broken spheres. Therefore, when the heating rate of calcination is 0.1 °C min−1, the spherical-like MgO nanostructures were obtained.
Fig. 2

a, b SEM images of MgO products at a heating rate of 5 °C min−1. c, d SEM images of

MgO products at a heating rate of 0.1 °C min−1

To further study the morphology of spherical-like MgO samples, the TEM images of spherical-like MgO nanostructures are shown in Fig. 3. Figure 3a indicated that the morphology of obtained MgO is spherical structure with an average diameter of ~ 200 nm, which is consistent with the results of SEM. The interparticle porosity of spherical-like MgO is clearly observed in the high-magnification TEM image (Fig. 3b). The TEM image also confirms that the porous of the sphere is constructed from small MgO nanoparticles. These conditions indicate that the obtained MgO spheres have a hollow mesoporous structure. It can be seen from Fig. 3c that lattice space of the sample is 0.24 nm, which corresponds to the (111) crystal planes of cubic MgO. Figure 3d shows the selected area electron diffraction (SAED) pattern of the corresponding sample, which has bright circled rings, indicating that spherical-like MgO has polycrystalline nature. The bright spots of the concentric rings indicate the (111), (222), (200), (220) planes of cubic MgO which are in good agreement with the MgO planes obtained by XRD.
Fig. 3

a, b TEM images of spherical-like MgO nanostructures. c HR-TEM image of spherical-like MgO nanostructures. d SAED pattern of spherical-like MgO nanostructures

The nitrogen adsorption–desorption isotherm was generally utilized to further analyze the pore structures of porous materials. Figure 4 shows the N2 adsorption–desorption isotherm and pore size distribution of the obtained MgO. The isotherm exhibits a typical Langmuir IV isotherm and has H3 hysteresis loop, indicating the presence of mesopores within the slit-like materials. According to the desorption data of the isotherm, the pore size distribution was calculated by BJH. It reveals a narrow pore size distribution in the mesoporous range for the as-prepared MgO sphere with the center at 4 nm (inset of Fig. 4), which further proves the pores are mesoporous. The specific surface area of the MgO sphere calculated from the nitrogen isotherm is 180 m2 g−1, which was almost higher than that of the being reported adsorbents (Table 1). Furthermore, the porous structures of the as-prepared MgO spheres supplied lots of surface active sites and greatly improved the adsorption rate and adsorption capacity.
Fig. 4

N2 adsorption–desorption isotherms of the MgO sample; insets show the corresponding pore size distributions

Table 1

Maximum adsorption capacities of different adsorbents for Pb(II) and Cd(II)


Surface area (m2·g−1)

Adsorption capacity (mg·g−1)

adsorbent dosage (g/l)

the concentrations of heavy metal











An et al. (2018)

Flower-like MgO






Cao et al. (2012)






Liao et al. (2018)

Micro/nanoscale magnesium silicate hollow spheres






Wang et al. (2010)

MgO nanoparticles






Xiong et al. (2015)

spherical MgO






This work

Lead (Pb(II)) and cadmium (Cd(II)) are two toxic heavy metal ions in drinking water resources, and how to efficiently remove them from water is a critical issue. In this experiment, we tested the adsorption ability of spherical MgO nanostructures as adsorbents for removing Pb(II) and Cd(II). Figure 5a shows the adsorption rates of Pb(II) and Cd(II) ions with an initial concentration of 100 mg L−1 by using the spherical-like MgO as absorbent versus contact time at room temperature, respectively. It is observed that the adsorption processes of Pb(II) and Cd(II) ions on MgO samples were very fast and there were no significant differences in the adsorption efficiency. Within the first 5 min, about 99% of Pb(II) and Cd(II) ions in the solution could be removed by MgO sample, and equilibriums were established after 10 min. Obviously, the time to reach adsorption equilibrium is much shorter than those of flower-like MgO, activated carbon (Cao et al. 2012). This rapid adsorption rate was due to the availability of abundant active sites on the surface of MgO spheres. As we all know, the concentration limits set by the World Health Organization for drinking water of Pb(II) and Cd(II) were 10 μg/l and 5 μg/l, respectively(Cao et al. 2012). In order to observe whether adsorption reached the limit, we extended the adsorption time. By extending the reaction time, it is possible to test whether the remaining Pb(II) and Cd(II) ion concentrations in the final solutions can reach the standard. The concentrations of the remaining Pb(II) and Cd(II) ions in the solutions are 99 μg/l and 97 μg/l, respectively. Although the concentrations do not reach the standard, they are lower than before. The adsorption isotherm was used to indicate the detailed relationship between the removal rate of a material and the concentration of the contaminant. We obtained adsorption isotherms with different concentrations ranging from 500 to 6000 mg L−1. The Langmuir adsorption model was utilized for the adsorption analysis. The formula was as follows:
$$ {q}_e={q}_mb{C}_e/\left(1+{K}_L{C}_e\right) $$
where Ce is the equilibrium concentration of metal ions (mg·L−1), qe is the amount of heavy metal ions adsorbed per unit weight of the adsorbent at equilibrium (mg·g−1), qm is the maximum adsorption capacity (mg/mg−1), and KL is the adsorption constant (L·mg−1) related to the energy of adsorption.
Fig. 5

a Adsorption rates and b adsorption isotherms of Pb(II) and Cd(II) with spherical MgO nanostructures as adsorbents

The qe of the adsorbent for Pb(II) and Cd(II) ions was calculated according to the equation:
$$ {q}_e=\left({C}_{\mathrm{o}}-{C}_e\right)V/m\kern0.5em $$
where Co is the initial concentrations of metal ions (mg·L−1), V is the volume of the solution (mL), and m is the weight of MgO(mg).

From Fig. 5b, we can see the experimental data agree well with the Langmuir adsorption isotherm, suggesting a surface adsorption process for heavy metal ions. The saturated adsorption capacities of the obtained MgO spheres for Pb(II) and Cd(II) are 5214 and 4187 mg g−1, respectively, which are much higher than those of nanomaterials, as shown in Table 1. The above results primarily showed that the spherical MgO could be used as a rapid and efficient adsorbent to remove metal ions for water treatment.

We also used the Freundlich isotherm to fit the experimental data. The formula was as follows:
$$ {q}_e=b{C_e}^{\raisebox{1ex}{$1$}\!\left/ \!\raisebox{-1ex}{$n$}\right.} $$
where b and 1/n are the adsorption constant, qe is the amount of heavy metal ions adsorbed per unit weight of the adsorbent at equilibrium (mg·g−1), Ce is the equilibrium concentration of metal ions (mg·L−1).
As shown in Fig. 6 and Tables 2 and 3, the Freundlich model has a poor fit for Cd(II) adsorption on MgO than Langmuir model. The R2 value obtained by Freundlich model is 0.969 and lower than that of Langmuir model, which was 0.976. It indicates that the adsorption isotherm data of MgO on Cd(II) ions do not fit with Freundlich model. However, the experimental data for Pb(II) adsorption on MgO were fitted well with Freundlich model and Langmuir model.”
Fig. 6

Fitting Freundlich adsorption isotherms of MgO for Cd(II) and Pb(II) ions

Table 2

Parameters by fitting of Langmuir and Freundlich adsorption models

Heavy metal ions



R2 (L)



R2 (F)



Table 3

Pseudo-second-order kinetic parameters for the adsorption of Pb(II), Cd(II) onto MgO


K2 (g/mg·min)

qe, exp. (mg·g−1)

qe, cal (mg·g−1)

R 2











For investigate the kinetics of the adsorption of Pb(II) and Cd(II) with spherical MgO as adsorbent, the pseudo-second-order kinetic model was utilized to test the experimental data. The equation is presented as
$$ \frac{t}{q_t}=\frac{1}{k_2{q}_e^2}+\frac{1}{q_e}\ t $$
where qt is the amount of heavy metal ions adsorbed per unit weight of the adsorbent at time t (mg·g−1), and K2 is the rate constant of adsorption (g/ mg/ min).
The pseudo-second-order rate is used widely in the study of adsorption kinetics. Figure 7 exhibits the linear plots of t/qt versus t of the Pb(II) and Cd(II), respectively. As shown in Table 3, all obtained R2 values were greater than 0.99. Moreover, the calculated q values (qe, cal) are close to the experimental q values (qe, exp). The kinetics data are consistent with the pseudo-second-order model, indicating that the specific surface area of MgO is high and the internal diffusion resistance is small, indicating that the rate determining step is the surface adsorption involving chemical adsorption.
Fig. 7

Pseudo-second-order kinetics model fitting of the plots for the adsorption of a Pb(II) and b Cd(II) onto MgO

We also used the pseudo-first-order kinetic model to fit the experimental data. The results are shown in Fig. SI and Table SI in the Supporting Information, which indicated that the data did not follow the first-order kinetics. The pseudo-first-order kinetic model was mainly utilized to describe the physical adsorption process.

The XRD was used to analyze the changes of powders after MgO adsorbing Pb(II) and Cd(II) ion solutions with initial concentration of 5000 mg/l. Figure 8a exhibits the XRD pattern of the sample that adsorbed for Pb(II); all the diffraction peaks can be assigned to Pb3(CO3)2(OH)2 (JCPDS 01-0687), PbO (JCPDS 72–0094) and Mg(OH)2 (JCPDS 84–2163), and the diffraction peaks for Cd(II) (as shown in Fig. 8b) can be indexed to Cd(OH)2 (JCPDS 84–1767), CdO (JCPDS 05–0640), Mg(OH)2 (JCPDS 84–2163), and CdCO3 (JCPDS 85–0989). For all we know, MgO nanoparticles can be hydrated to form Mg(OH)2 in water (Cao et al. 2012), and our experiment results also demonstrate the consideration. Therefore, the Mg(OH)2 phase was produced by hydration of MgO in aqueous solution. Furthermore, the generated part of Mg(OH)2 would be dissociated to form OH ions close to the surface of MgO nanoparticles. Therefore, Pb(II) and Cd(II) in the solutions could combine with the produced OH ions forming insoluble hydroxides (Pb(OH)2, Cd(OH)2). Since the unstable formation of Pb(OH)2, it could react with dissolved CO2 to form Pb3(CO3)2(OH)2 (Bian et al. 2010). The existence of CdO and PbO phases could be attributed to the cation exchange between Mg(II) and Cd(II) or Pb(II) ions, suggesting that Cd(II) or Pb(II) ions replace Mg(II) in the lattice (Cao et al. 2012). Actually, it is a solid–liquid interfacial reaction that can combine with metal ions (Cd2+ and Pb2+) to form metal oxides (CdO and PbO). The reactions are as follows:
$$ \mathrm{Mg}\mathrm{O}+{\mathrm{H}}_2\mathrm{O}\kern0.5em \to \mathrm{Mg}{\left(\mathrm{OH}\right)}_2 $$
$$ {\mathrm{Cd}}^{2+}+\mathrm{Mg}{\left(\mathrm{OH}\right)}_2\kern0.5em \to \kern0.5em \mathrm{Cd}{\left(\mathrm{OH}\right)}_2+{\mathrm{Mg}}^{2+} $$
$$ {\mathrm{Cd}}^{2+}+{\mathrm{CO}}_2+{\mathrm{H}}_2\mathrm{O}\to \kern0.5em \mathrm{Cd}{\mathrm{CO}}_3+2{\mathrm{H}}^{+} $$
$$ 3{\mathrm{Pb}}^{2+}+\mathrm{Mg}{\left(\mathrm{OH}\right)}_2+2{\mathrm{CO}}_2+2{\mathrm{H}}_2\mathrm{O}\to {\mathrm{Pb}}_3\ {\left({\left(\mathrm{CO}\right)}_3\right)}_2{\left(\mathrm{OH}\right)}_2+{\mathrm{Mg}}^{2+}+4{\mathrm{H}}^{+} $$
$$ \mathrm{MgO}+{\mathrm{Pb}}^{2+}\to \mathrm{PbO}+{\mathrm{Mg}}^{2+} $$
$$ \mathrm{MgO}+{\mathrm{Cd}}^{2+}\to \mathrm{CdO}+{\mathrm{Mg}}^{2+} $$
Fig. 8

XRD patterns of MgO nanostructures after adsorbing Pb(II) (a) and Cd(II) (b)

To further reveal the removal mechanism of Pb(II) and Cd(II) by MgO sample, the FTIR was used to identify the functional groups of MgO before and after adsorbing Pb(II) and Cd(II) ion solutions with initial concentration of 5000 mg/l. Figure 9a shows the mixed spectrum of MgO and Mg(OH)2, the peak at 3430 cm−1 corresponds to the –OH stretching mode of hydroxyl groups in water molecule, and the sharp strong peak at 3701 cm−1 is attributed to the vibration of Mg(OH)2 (Liu et al. 2015). The band at 1645 cm−1 is commonly assigned to the bending vibration of –OH (Liu et al. 2015). Band at 1447 cm−1 is the stretching vibration of CO32− (Chowdhury et al. 2016). The peaks at 857 and 571 cm−1 are due to the IR characteristic bands of MgO (Chowdhury et al. 2016). Figure 9b shows the spectrum of MgO after adsorbing Pb(II), the peak at 3701 cm−1 is almost disappeared, suggesting that the hydroxyl groups of Mg(OH)2 may participate in the adsorption process. In addition, band at 1043 cm−1 is the stretching vibration of CO32−; the possible cause of above conditions may be due to the formation of Pb3(CO3)2(OH)2 on the surface of MgO. Peaks at about 687 and 522 cm−1 are due to the characteristic adsorption peak of the PbO (Senvaitiene et al. 2007), while the peak at 571 cm−1 for MgO is not visible, which may be covered by the PbO. The FTIR spectrum of MgO after adsorbing Cd(II) ion solution is shown in Fig. 9c. It is obviously observed that the peak of Mg(OH)2 has moved to 3742, and the percentage transmittance is increased. These phenomena might be caused by the vibration bands of Cd(OH)2 formed on the surface of MgO after adsorbing Cd(II) ion solution. Moreover, the peak at 3600 cm−1 is attributed to the Cd(OH)2 (Liu et al. 2015). The peak at 1405 cm−1 is the stretching vibration of CO32− (Cui et al. 2018; Janet et al. 2007). Band at 716 cm−1 corresponds to the presence of CdCO3 vibration, while the peak at 440 cm−1 is the stretching vibration band of CdO (Marzouk et al. 2014). The same as the adsorption of Pb(II) ion, the peak at 571 cm−1 for MgO is not visible, which means that it is covered by the CdO. Consequently, it can be concluded from FTIR analysis that hydroxyl functional groups play a major role in the removal of Pb(II) and Cd(II) by using obtained MgO.
Fig. 9

FTIR spectra of a MgO and Mg(OH)2. b MgO after adsorbing Pb(II). c MgO after adsorbing Cd(II)

Therefore, the mechanism of removing Pb(II) and Cd(II) ions from aqueous solution by MgO was considered to be cation exchange between Mg(II) and heavy metal ions on spherical MgO. New substances were formed after adsorption. The adsorption of MgO on Pb(II) and Cd(II) ions was chemical adsorption, which leads to the change of the morphology of MgO. Therefore, it is difficult to regenerate.


In summary, we demonstrated that spherical-like MgO nanostructures were facile to synthesize by spray-drying method. They exhibit high surface area and improved adsorption properties in the removal of heavy metal ions and show maximum capacities up to 5214 mg/g and 4187 mg/g for Pb(II) and Cd(II), respectively. These values are much higher than those reported MgO-based adsorbents. The adsorption kinetics and isotherm data agree well with pseudo-second-order and Langmuir models, respectively, indicating the monolayer chemisorption of heavy metal ions. Moreover, in less than 10 min, Pb(II) and Cd(II) ions in solution can be almost removed, which means that the spherical-like MgO possesses a high adsorption rate. The mechanism of removing Pb(II) and Cd(II) was ion exchange between Mg and heavy metal ions. Because of the advantages such as high adsorption capacity, high surface areas, and low cost, the spherical-like MgO nanostructures would be a potential adsorbent for removing heavy metal ions from wastewater.


Funding information

This work was supported by National Natural Science Foundation of China (51664023, 51561007), Natural Science Foundation of Jiangxi Province (20181BBE58001, GJJ181510), and Natural Science Foundation of Ganzhou City (20170215, 20180143).

Supplementary material

11356_2019_5277_MOESM1_ESM.docx (78 kb)
ESM 1 (DOCX 77 kb)


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Copyright information

© Springer-Verlag GmbH Germany, part of Springer Nature 2019

Authors and Affiliations

  1. 1.School of Materials Science and EngineeringJiangxi University of Science and TechnologyGanzhouChina
  2. 2.Institute of Research and EngineeringJiangxi University of Science and TechnologyGanzhouChina

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