A waxberry-like SiO₂@MnSiO₃ core–shell nanocomposite synthesized via a simple solvothermal self-template method and its potential in catalytic degradation and heavy metal ion removal

Abstract

A novel waxberry-like SiO₂@MnSiO₃ core–shell nanocomposite was facilely fabricated via the simple one-step thermal treatment of SiO₂ nanospheres, MnCl₂·4H₂O, ethylenediamine (EDA), and ethylene glycol (EG). Through an intensive investigation of the effects of Si/Mn molar ratio and reaction time on the grain growth characteristics, a self-template growth mechanism of SiO₂@MnSiO₃ was proposed. The self-template silica nanospheres released silicate anions slowly from their surfaces by alkali etching in the presence of EDA, and a fast precipitation reaction between Mn²⁺ cations and silicate anions occurred within the interfacial regions, eventually leading to the formation of a MnSiO₃ shell on the surfaces of silica nanospheres. A well-defined waxberry-like SiO₂@MnSiO₃ nanostructure was obtained with a Si/Mn molar ratio of 5 : 1 and a reaction time of 10 h according to our experiments. Interestingly, this SiO₂@MnSiO₃ exhibited a high catalytic activity for oxidative degradation of methylene blue (MB); more than 93% of MB could be decomposed within 40 min. Moreover, it could also act as a potential adsorbent for efficient removal of Pb²⁺ ions from aqueous solution. The Pb²⁺ adsorption capacity was up to 50.5 mg g⁻¹, which was significantly higher than those found for many other conventional adsorbents. Overall, this work not only provides a new insight into the fabrication of silica-supported MnSiO₃ nanocomposites but also demonstrates their excellent performance in heterogeneous catalysis and adsorption.

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1. Introduction

Inorganic manganese-based materials, such as MnO, Mn₃O₄, MnO₂, MnOOH, MnSiO₃, and MnTiO₃, are of considerable current interest owing to their great potential applications in heterogeneous catalysis and adsorption.¹ The catalytic activities of manganese-based materials are mainly attributed to the remarkably versatile redox chemistry of Mn with formal oxidation states II–V and VII,² while their adsorptive properties are in large part due to the fact that the surface coordinatively unsaturated Mn sites are readily converted into Mn–OH groups in aqueous solution, providing manganese-based materials with considerable ability to capture a variety of organic or inorganic species.^3,4 As far, great efforts have been made to enhance the catalytic or adsorptive performance of manganese-based materials through regulation of their particle size, morphology, and porosity.^5–8 There is now a high degree of agreement among scientific experts that the nanostructured manganese-based materials should be regarded as preferred alternatives because they always show more excellent performance characteristics when compared with the bulk ones.⁹ However, the unsupported manganese-based nanoparticles frequently undergo conglomeration due to their high surface energy, resulting in loss of catalytic activity or reduction in adsorption capacity.¹⁰ The use of solid materials to synthesize supported manganese-based nanocomposites seems to be of practical significance because it can inhibit the aggregation of neighboring manganese-based nanoparticles and expose more active sites for catalysis or adsorption.¹¹

Among various supports available, silica has been considered a particularly appropriate candidate because its morphology and size can be easily and precisely controlled.¹² Moreover, silica has a good thermal stability and mechanical strength.¹³ There are three general strategies for fabrication of silica-supported manganese-based nanocomposites.¹⁴ The most popular method is the co-hydrolysis/condensation, which involves the direct addition of manganese precursors to the synthesis solution of silica.¹⁵ The method is simple, efficient, and requires less time than other synthetic approaches. However, this approach has its associated drawbacks. The presence of manganese precursors in solution may interfere with the polymerization chemistry of silica, often resulting in product with undesirable properties including decreased mechanical strength and less well defined morphology.¹⁴ Moreover, a large part of the Mn sites will be inactive because they are buried inside the silica matrix.¹⁴ An alternative to the co-hydrolysis/condensation method is wetness impregnation.¹⁶ Generally, this method includes the following steps: (i) preparation of silica colloids; (ii) impregnation of silica colloids with manganese precursor solution; and (iii) the impregnated solid is oven dried and subsequently calcined. Although the wetness impregnation method is experimentally straightforward, it suffers from a poor control over the growth of manganese-based nanoparticles on silica, affecting their dispersion and size with increasing loading amount.¹⁴ Microemulsion method is a new route for synthesizing silica-supported manganese-based nanocomposites that has been developed in recent years.^17,18 In this method, the first step is the formation of a water-in-oil microemulsion containing manganese precursor (or manganese-based nanoparticles) and silica precursor, followed by the hydrolysis/condensation process. Silica-supported manganese-based nanocomposites using this method have been reported to have a more controllable composition and particle size compared with traditional methods.¹⁹ Nevertheless, like co-hydrolysis/condensation method, the microemulsion method also causes a significant amount of Mn sites to be covered by silica.²⁰ Moreover, the method uses excessive surfactant (emulsifier) molecules and requires a substantial number of washing steps, which seems to be costly and tedious.

Our interest is to explore a facile synthesis method without the drawbacks mentioned above to develop silica-supported manganese-based nanocomposite for both catalysis and adsorption applications. Through our ongoing efforts, herein we proposed a simple solvothermal self-template strategy for fabrication of a waxberry-like SiO₂@MnSiO₃ core–shell nanocomposite, using monodisperse silica nanospheres, manganese chlorite (MnCl₂·4H₂O), ethylenediamine (EDA), and ethylene glycol (EG) as self-template, manganese source, organic alkaline, and solvent, respectively. To the best of our knowledge, this is the first work that examines self-template method for fabrication of SiO₂@MnSiO₃. The synthesis process is schematically illustrated in Fig. 1. Briefly, the EDA could stabilize Mn²⁺ ions in EG by chelating interactions, while silica nanospheres released soluble silicate by alkali etching. As a result, a fast precipitation reaction between Mn²⁺ and silicate species occurred within the interfacial regions and eventually led to the formation of MnSiO₃ shell on the surfaces of silica nanospheres.

Fig. 1 A schematic illustration of the formation of the waxberry-like SiO₂@MnSiO₃ core–shell nanostructure in the synthetic process.

The desired SiO₂@MnSiO₃ might have the following notable features: (i) both SiO₂ core and MnSiO₃ shell are environmentally friendly materials with good thermal stability and mechanical strength;²¹ (ii) the MnSiO₃ nanograins are highly dispersed on the surfaces of SiO₂ and the final core–shell architecture has a regular waxberry-like morphology, ensuring the maximum exposure of active Mn sites; and (iii) the SiO₂@MnSiO₃ has a perfect spherical shape with narrow size distribution, make it ideal candidate as catalyst or adsorbent. Considering the above advantages and in continuation of our research work on designing heterogeneous catalytic systems and highly efficient adsorbents for wastewater purification, we studied the catalytic property of SiO₂@MnSiO₃ by employing it as Fenton-like catalyst for organic dye (i.e., methylene blue) degradation and tentatively used it as an adsorbent for removal of heavy metal ion (i.e., Pb²⁺).

2. Experimental

2.1 Chemicals and reagents

Ethylene glycol (EG), ethylenediamine (EDA), manganese chloride tetrahydrate (MnCl₂·4H₂O), tetraethyl orthosilicate (TEOS), methylene blue (MB), hydrogen peroxide (H₂O₂, 30%, w/w), and lead(II) nitrate (Pb(NO₃)₂) were all obtained from Sinopharm Chemical Reagent Co., Ltd. (Shanghai, China). Distilled water was used throughout the experiments for solution preparation.

2.2 Preparation of SiO₂@MnSiO₃

The SiO₂ nanospheres were synthesized via the Stöber method.²² Typically, 61.75 mL absolute ethanol and 4.5 mL TEOS were added into 24.75 mL deionized water, followed by vigorously stirring for 10 min. Then, 9 mL ammonium hydroxide solution (∼25 wt%) was added. The resultant mixture was continuously stirred for 2 h to form a suspension. The solid products (SiO₂ nanospheres) were collected by centrifugation and washed repeatedly with water/ethanol. Afterwards, the clean SiO₂ nanospheres were dispersed in EG for subsequent use. The concentration of SiO₂ nanospheres was about 89.6 mg per milliliter.

The formation of MnSiO₃ shell on SiO₂ core was achieved via a solvothermal process. Briefly, 98.8 mg MnCl₂·4H₂O was dissolved in a certain amount of EG. Then, a certain amount of SiO₂ suspension and 0.034 mL EDA were added. Total volume of the mixture was fixed at 20 mL. In order to investigate the effect of Si/Mn molar ratio on the grain growth characteristics, five different Si/Mn molar ratios, 2 : 1, 3 : 1, 4 : 1, 5 : 1, and 6 : 1, were tested. After stirring for 10 min, the mixture was sealed in a 50 mL Teflon-lined autoclave and maintained at 200 °C for 1–20 h. The solid product (i.e., SiO₂@MnSiO₃) was collected by centrifugation, washed repeatedly with water/ethanol, and then dried in a vacuum oven at 60 °C overnight. To completely remove the EG and/or EDA residue, the SiO₂@MnSiO₃ nanoparticles were treated at 400 °C for 2 h with a heating rate of 2.5 °C min⁻¹. The obtained SiO₂@MnSiO₃ products were denoted as SMS-x-y, where the x and y represented solvothermal treatment time and Si/Mn molar ratio, respectively.

2.3 Characterization of materials

The morphological analyses were examined on a Philips CM-12 transmission electron microscope (TEM, 200 kV) and a Hitach SU8010 field emission scanning electron microscope (SEM, 15 kV). The powder X-ray diffraction (XRD) measurements were recorded on an X-ray diffractometer (X'Pert Pro DY2189, PANaytical, B. V., the Netherlands) using Cu-Kα radiation with scattering angles. The N₂ adsorption/desorption analysis was performed on a Micromeritics ASAP2020 surface area analyzer at 77 K.

2.4 Catalytic degradation procedure

The catalytic property of SMS-10-5 samples were evaluated by the oxidative degradation of MB via a Fenton-like chemical reaction. Typically, 10 mL of MB dye solution (50 mg L⁻¹) containing 1.5 mL of 30 wt% H₂O₂ was added into a vial that included 5 mg of catalysts, and the mixture was stirred at 60 °C. At predetermined time intervals, the catalyst was separated immediately from the solution by filtration. The concentrations of MB in the course of degradation were measured at the maximum absorption wavelength (665 nm) by using the UV-vis spectrophotometer, and the decoloring degree (DD) of MB was calculated by the following equation:where C₀ and C_t were concentrations of MB (mg L⁻¹) at initial and predetermined time, respectively.

2.5 Adsorption procedure

In a typical procedure to obtain the adsorption isotherm, 5 mg SMS-10-5 were mixed with 10 mL Pb²⁺ aqueous solution (pH 6.0) of different concentrations (10–100 mg L⁻¹). The resulting mixture was shaken with speed of 260 rpm in a shaker at 303.15 K for 3 hours to ensure adsorption equilibrium. The Pb²⁺-adsorbed SMS-10-5 nanoparticles were then separated by centrifugation. The adsorption amount of Pb²⁺ was determined by UV-vis spectrophotometry using xylenol orange as color reagent and phenanthroline as masking agent, and its value was calculated by the following equation:where q_e is the amount of Pb²⁺ adsorbed at equilibrium (mg g⁻¹), C₀ and C_e are the initial and equilibrium concentrations of Pb²⁺ (mg L⁻¹), m is the mass of SMS-10-5 (g), and V is the volume of solution (L).

In order to determine the adsorption rate and kinetic characteristics, experiments were conducted by varying the contact time from 2 to 240 min at initial Pb²⁺ concentrations of 80 mg L⁻¹.

3. Results and discussion

3.1 Characterization of waxberry-like SiO₂@MnSiO₃ core–shell structure

The SiO₂ nanospheres were prepared via the classical Stöber method. As shown in Fig. 2a, the SiO₂ nanospheres exhibit monodisperse, uniform, and spherical morphology with an average diameter about 440 nm. Moreover, the surface of the SiO₂ nanospheres is relatively smooth because of the amorphous nature of silica.²³ After solvothermal treatment of SiO₂ nanospheres in EG solution containing EDA and Mn²⁺ ions at 200 °C for 10 h with Si/Mn molar ratio of 5 : 1, the as-fabricated SiO₂@MnSiO₃ (denoted as SMS-10-5) nanoparticles are found to be composed of monodisperse waxberry-like core–shell nanospheres with an average diameter about 460 nm (Fig. 2b). Moreover, the outer shell contacts with the silica core closely (Fig. 2b). A closer look shows that the shell of SMS-10-5 is hierarchical and assembled by tiny nanoparticles, and the thickness of the shell is about 20 nm (Fig. 2c and d). It was expected that this unique structure might have a relatively high surface area due to the hierarchical characteristic and could provide high catalytic and adsorptive efficiencies because of the easy access to its active sites.

Fig. 2 SEM images of SiO₂ (a) and SMS-10-5 (b and c); TEM image of SMS-10-5 (d).

Fig. 3a shows the XRD patterns of the self-template SiO₂ nanospheres and SMS-10-5. The diffraction pattern of the SiO₂ shows a broad peak around 23°, indicating its amorphous nature.²⁴ After solvothermal treatment, the diffraction pattern of the resulting material (SMS-10-5) shows a reflection characteristic of MnSiO₃ (JCPDS Card no. 12-0181) in addition to the SiO₂ reflection.²⁵ This result indicates that the MnSiO₃ shell has been successfully formed onto the SiO₂ core. Furthermore, the chemical state of manganese element in SMS-10-5 is analyzed by XPS (Fig. 3b). The result shows that the Mn 2p_1/2 and 2p_3/2 signals are situated at 656.1 and 642.8 eV, respectively. The result indicates the valence state of Mn is divalent (+2) in the shell of SMS-10-5, which is well consistent with the literature value.²⁶

Fig. 3 XRD patterns of SiO₂ and SMS-10-5 (a), XPS spectrum of SMS-10-5 (b), EDS mapping of SMS-10-5 showing the distribution of Mn, O and Si elements (c), and N₂ adsorption/desorption isotherms (the inset showing the pore size distribution of SMS-10-5) (d).

Elemental mapping based on energy dispersive spectroscopy (EDS) of SMS-10-5 verifies that Mn, Si, and O are uniformly distributed in the hierarchical nanostructures, indicating the MnSiO₃ is uniformly deposited on the surface of the SiO₂ core (Fig. 3c). Moreover, according to the EDS result, the mass content of manganese in the SMS-10-5 is found to be about 7 wt%.

The specific surface area and hierarchical porous nature of the as-prepared waxberry-like SMS-10-5 nanostructures were further investigated by nitrogen (N₂) adsorption–desorption measurement, as indicated in Fig. 3d. For a comparison, the template SiO₂ nanospheres were also measured and the corresponding curve is depicted in Fig. 3d. It can be seen that the curve of SMS-10-5 exhibits an obvious hysteresis loop in the high relative pressure region (P/P₀ = 0.6–0.9), exhibiting the characteristics of a porous material, which is consistent with the above microscopy observations. The porous structure is further supported by the Barrett–Joyner–Halenda (BJH) pore size distribution shown in Fig. 3d. Apparently, the pore size range of SMS-10-5 is distributed between 2 and 20 nm (Fig. 3d). The Brunauer–Emmett–Teller (BET) surface area of the waxberry-like SMS-10-5 nanostructures is thus calculated to be around 38.1 m² g⁻¹, which is significantly larger than the BET surface area of the SiO₂ nanospheres with a value of 15.9 m² g⁻¹. With a large surface area and porous structure, the as-synthesized waxberry-like SMS-10-5 nanocomposite should have potential application in heterogeneous catalysis and adsorption, since they may facilitate the molecule- or ion-transfer at the solid–liquid interface.

3.2 Growth mechanism of waxberry-like SiO₂@MnSiO₃ core–shell structure

3.2.1 Theoretical analysis. As depicted in Experimental section, the SiO₂@MnSiO₃ samples were synthesized via one-step thermal treatment of SiO₂ nanospheres, MnCl₂·4H₂O, EG, and EDA. It is well known that EDA can form a stable complex with Mn²⁺ ion according to the following reaction:²⁷

Mn²⁺ + NH₂(CH₂)₂NH₂ ↔ [Mn(NH₂(CH₂)₂NH₂)₂]²⁺ (3)

Besides, as an organic alkaline, EDA is also likely to hydrolyze and release OH⁻ ion. The reaction can be expressed as follows:²⁸

NH₂(CH₂)₂NH₂ + 2H₂O ↔ [NH₂(CH₂)₂NH₂]²⁺ + 2OH⁻ (4)

where the H₂O molecule might be the chemically combined water of MnCl₂·4H₂O or result from the self-condensation reaction of EG molecules.^29,30

In theory, it seems that the Mn²⁺ should also react with OH⁻ ion to form Mn(OH)₂ precipitate as follows:³¹

Mn²⁺ + 2OH⁻ → Mn(OH)₂↓ (5)

Actually, although the data is not available, it might be expected that EDA can form stronger binding force with Mn²⁺ than OH⁻ ion. Containing two nitrogen atoms that can coordinate to a metal center, EDA is a multidentate ligand.²⁹ It is a general observation that chelated complexes of multidentate ligands are more thermodynamically stable than those of the same metal ion with monodentate ligands.³² Moreover, in our synthesis system, the concentration of EDA should be much higher than that of OH⁻ ions. Hence, the formation of the stable complex [Mn(NH₂(CH₂)₂NH₂)₂]²⁺ should effectively suppress the precipitation reaction between Mn²⁺ and OH⁻. This is supported by the phenomenon that no precipitates but only transparent solution was obtained when the reaction was solvothermally treated under 200 °C for 10 h without the addition of SiO₂ (not shown).

On the other hand, amorphous silica can be etched away by OH⁻ as follows:³³

SiO₂ + 2OH⁻ ↔ SiO₃²⁻ + H₂O (6)

Because of the fairly low concentration of OH⁻ in our synthesis system, it might be expected that the reactivity of silica etching should be very low. To clarify this, a control experiment, i.e., thermal treatment of SiO₂ spheres, EG, and EDA under 200 °C for 10 h without the addition of MnCl₂·4H₂O, was conducted. As a result, we detected no recognizable differences between the raw SiO₂ nanospheres and the ones obtained from the SiO₂/EG/EDA system (not shown). The result clearly demonstrated that the silica etching only had a very low efficiency in absence of MnCl₂·4H₂O.

However, when MnCl₂·4H₂O is introduced into the SiO₂/EG/EDA system, a thermodynamically favorable precipitation reaction between Mn²⁺ and silicate species may occur within the interfacial region as follows:²⁵

Mn²⁺ + SiO₃²⁻ → MnSiO₃↓ (7)

Driven by the reaction shown in eqn (7), the SiO₂ nanospheres will be gradually dissolved, and MnSiO₃ nanograins will be grown readily around the surface of the spherical SiO₂ core to form a MnSiO₃ shell.

3.2.2 Si/Mn and time dependent experiments and formation mechanism. To better understand the formation mechanism of the as-synthesized waxberry-like nanostructures, batch experiments dependent on the molar ratios of Si/Mn were firstly carried out. The morphologies and structures of the resulting products were characterized. Fig. 4 presents SEM and TEM images of the resulting products synthesized with different molar ratios of Si/Mn. As the molar ratios of Si and Mn increase from 2 : 1 to 5 : 1, all the resulting products (SMS-10-2, SMS-10-3, SMS-10-4, and SMS-10-5) roughly exhibit waxberry-like morphologies (Fig. 4a–d and f–h). After careful observation, it might be noticed that a small amount of MnSiO₃ nanoparticles stand outside the as-formed core–shell nanostructures in the first three samples (Fig. 4a–c, f and g), while no redundant MnSiO₃ nanoparticles can be observed in SMS-10-5 sample (Fig. 4d and h). Apparently, the molar ratio of Si/Mn is key to the formation of a well-defined SiO₂@MnSiO₃ core–shell nanostructure. Under identical reaction conditions (both amounts of Mn²⁺ and OH⁻ ions are constant), a lower SiO₂ concentration (corresponding to lower Si/Mn value) means that the average SiO₂ nanosphere can be attached by more OH⁻ ions. In cases of SMS-10-2, SMS-10-3, and SMS-10-4, the surface area of SiO₂ nanosphere might be not large enough to accommodate all the newly generated MnSiO₃ nanograins, resulting in that a part of MnSiO₃ nanograins would be isolated from the shells. The lower the concentration of SiO₂ nanospheres is, the more will the isolated MnSiO₃ nanograins be formed. Conversely, as the concentration of SiO₂ nanospheres enhances, the amount of MnSiO₃ nanograins apportioned to each SiO₂ nanosphere will decrease, and thus there will be a better control over the growth of MnSiO₃ nanograins on SiO₂ nanospheres. In particular, when the molar ratio of Si/Mn is up to 5 : 1, the resultant product (SMS-10-5) shows a well-defined waxberry-like nanostructure, with no evidence of phase segregation at nanoscale (Fig. 4d and h). When the molar ratio of Si and Mn is further increased to 6 : 1, the resultant product (SMS-10-6) also shows a well-defined nanostructure. However, the thickness of the shell of SMS-10-6 is quite thin (Fig. 4e and i), which might be less advantageous for its application performance. Overall, a suitable concentration of SiO₂ nanospheres is needed to prepare SiO₂@MnSiO₃ with desired core–shell structure. The detailed Si/Mn molar ratio-dependent evolution of SiO₂@MnSiO₃ nanostructure is depicted in Fig. 4j.

Fig. 4 SEM images of SMS-10-2 (a), SMS-10-3 (b), SMS-10-4 (c), SMS-10-5 (d), and SMS-10-6 (e); TEM images of SMS-10-3 (f), SMS-10-4 (g), SMS-10-5 (h), and SMS-10-6 (i); the detailed Si/Mn molar ratio-dependent evolution of SiO₂@MnSiO₃ nanostructure (j).

To further study the formation process of SiO₂@MnSiO₃ core–shell nanostructure, time-dependent experiments were also performed, fixing the Si/Mn molar ratio at 5 : 1. Representative SEM images of the products (SMS-1-5, SMS-5-5, SMS-10-5, and SMS-20-5) at different reaction times are presented in Fig. 5a–d. As shown in Fig. 5a, at the early stage of the reaction (1 h), the originally smooth surface of the SiO₂ nanospheres becomes slightly rougher because of the formation of MnSiO₃ nanograins. With the reaction proceeding (5 h), the shells become more rough, indicating that more MnSiO₃ grow onto the SiO₂ surface (Fig. 5b). When the reaction time is prolonged to 10 h, a well-defined waxberry-like nanostructure is obtained (Fig. 5c). Even the reaction time extends to 20 h, the above waxberry-like nanostructures remain almost unchanged (Fig. 5d). These results indicate that the solvothermal reaction can be accomplished within 10 h. Also, it can be inferred that the MnSiO₃ shell enhances the stability of SiO₂ core against the subsequent etching by OH⁻.

Fig. 5 SEM and TEM images of SMS-1-5 (a), SMS-5-5 (b), SMS-10-5 (c), and SMS-20-5 (d).

On the basis of the above-obtained results, a self-template formation mechanism was proposed (Fig. 1). When solvothermally treated in Mn²⁺ and EDA solution, the surface of SiO₂ nanospheres is gradually dissolved in the form of silicate anions. Driven by the interfacial reaction between Mn²⁺ cations and the silicate anions, MnSiO₃ nanograins are grown readily around the surface of SiO₂ to form MnSiO₃ shell. Herein, the SiO₂ core serves not only as the precursor for the MnSiO₃ shell but also as a sacrificial template for the core–shell structure. When a high concentration of SiO₂ nanospheres is used (e.g., Si/Mn > 5 : 1), the SiO₂ can provide a surface that is large enough to hold MnSiO₃ nanograins, possibly resulting in a well-defined core–shell structure. On the other hand, when the concentration of SiO₂ nanospheres is low (e.g., Si/Mn < 5 : 1), the surface of SiO₂ nanosphere is insufficient to accommodate all the MnSiO₃ nanograins, resulting in a phase segregation. With the reaction proceeding, more and more MnSiO₃ nanograins grow onto the surface of SiO₂ nanospheres, which not only results in the formation of MnSiO₃ shells but also inhibit the etching of SiO₂ nanospheres. After a specified time (i.e., 10 h for the specimen with Si/Mn of 5 : 1), a waxberry-like SiO₂@MnSiO₃ core–shell can be obtained.

3.3 Catalytic performance

Organic compounds (e.g., dyes) are one of the most important groups of pollutants present in water. The development of powerful and practical treatments of organic compounds in wastewater has attracted world-wide attention in recent years.³⁴ Manganese-based nanomaterials of different valence have been reported to function as effective catalysts toward oxidative degradation of organic pollutants.² A notable prior work was reported by Tušar et al. who showed for the first time that manganese silicate nanoparticles could act as a superior Fenton-type catalyst toward oxidative degradation of organic pollutant.³⁵ Due to the abundant and highly accessible active sites present in MnSiO₃ shell, we expected that the SMS-10-5 from our experiment would be useful in Fenton-type catalysis.

The catalytic activity of SMS-10-5 was investigated by using it as the catalyst for oxidative degradation of organic pollutant. Methylene blue (MB), as a typical industrial pollutant, was chosen as a model. As the degradation of MB proceeds, the characteristic absorption of MB at 665 nm gradually weakens, which was why it was chosen for monitoring the catalytic process of SMS-10-5. The UV-vis absorption spectra of MB were measured as a function of the catalytic reaction time with SMS-10-5 as catalyst in the presence of H₂O₂ (Fig. 6a). The spectrum at t = 0 was taken on the starting solution of MB with a concentration of 50 mg L⁻¹. Obviously, the MB absorption peaks plummet rapidly (Fig. 6a). Within 40 min, the peaks at 665 nm are roughly reduced by an order of magnitude. As the reaction proceeded, the color of the solution changed from dark blue to near colorless, suggesting the highly efficient degradation of MB molecules. Moreover, it is confirmed that in the absence of any catalyst (only MB + H₂O₂), no obvious dye decoloration can be observed even after 90 min (Fig. 6b). With the SMS-10-5 but no H₂O₂ (only MB + SMS-10-5), the degree of decoloration reaches about 17% within 30 min and then tends to be saturated (Fig. 6b). The decoloration of the dye solution can be ascribed to adsorption of the dye molecules on the SMS-10-5 catalyst. When both SMS-10-5 catalyst and H₂O₂ oxidant are added to the MB solution, obvious decoloration occurs. Remarkably, the use of SMS-10-5 as a catalyst allows the degree of decoloration to reach 93% in only 40 min (Fig. 6b). The saturated degree of decoloration is as high as 96% (Fig. 6b). These results indicate that the decoloration of MB molecules is caused by H₂O₂-induced oxidation, catalyzed by SMS-10-5.

Fig. 6 Normalized UV-vis spectra of MB vs. reaction time, where the inset shows the photographs of MB solutions at different time interval (a); and normalized concentrations of MB vs. reaction time (b).

To describe MB-decolorizing processes more clearly, the kinetic data were fitted to the following kinetic equation:³⁶

Then,

where C_t and k are the MB concentration at time t and the rate constant, respectively.

The fitting results are summarized in Table 1. It can found that, the initial rate constant (k, min⁻¹) of the reaction was calculated as 0.1306 min⁻¹. Compared with the reported homogeneous and heterogeneous Fenton systems,^36,37 these SMS-10-5-based systems also show a higher initial rate constant, indicating that the catalytic oxidation process progresses more quickly.

Table 1 Comparison of the maximum adsorption capacities of MB between various adsorbents

Adsorbent	Maximum adsorption capacity (mg g^−1)	Adsorption temperature (K)	Ref.
Carbon nanotubes	30.32	298	46
Pyrolyzed coffee residues and clay	19.5	303	47
Sludge-derived biochar	30.88	298	48
MnO2-loaded resin	80.64	298	49
Al2O3-supported iron oxide	29.0	318	50
Activated phosphate rock	15.47	313	51
Phosphate rock	12.78	313	51
Phosphatic clay	37.2	298	52
SMS-10-5	50.5	303.15	This work

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The catalytic performance was described only by the degradation degree of MB might be unconvincing and improper because the weight of catalyst, the volume of H₂O₂ solution used, the initial concentration of MB, and the volume of MB solution were all responsible for the degradation of MB. Recently, the catalytic performance of catalyst has been proposed to be estimated by the equation as follows:³⁶

where Q (g L⁻¹ g⁻¹) is the consumption of MB caused by 1 g of catalyst and 1 L of 30 wt% H₂O₂ solution; V (mL) and V′ (mL) are the volume of the MB solution and the volume of 30 wt% H₂O₂ solution, respectively; m (g) is the mass of the catalyst; C₀ (g L⁻¹) and C (g L⁻¹) are the concentrations of MB at the beginning and ending of degradation, respectively. The Q value in this study is about 63.7 g L⁻¹ g⁻¹, which is significantly larger than 8.3 g L⁻¹ g⁻¹ of β-MnO₂ nanorods, 15.0 g L⁻¹ g⁻¹ of mesoporous manganese ferrite nanocomposite, and 36.7 g L⁻¹ g⁻¹ of magnetic Fe₂MnO₄ activated carbons.^38–40 The excellent catalytic performance of the as-prepared SMS-10-5 might be related to its unique waxberry-like nanostructure.

3.4 Adsorption properties

The discharge of toxic heavy metal ions into water is another serious pollution environmental problem. Different from organic pollutants, heavy metal ions are not biodegradable and therefore persistent. They tend to accumulate in the living organisms and enter the food chains through various pathways, which bring many detrimental effects on environment and human health.⁴¹ Among all of the metal ion removal methods, the adsorption technique has considered as the most extensively adopted method because its cost-effectiveness, simple operation, and high efficiency.^41–45 Recent studies have noted that Mn-based materials can act as promising adsorbents for heavy metal ion removal.^3,4 This is partly because the surface coordinatively unsaturated Mn sites are readily converted into Mn–OH groups in an aqueous solution, endowing Mn-based materials with considerable ability to capture a variety of heavy metal ions.^3,4 Therefore, here we investigated the adsorption behavior of SMS-10-5 for the treatment of lead aqueous solutions.

The isothermal experiments of Pb²⁺ adsorption onto SMS-10-5 was conducted, and the results are shown in Fig. 7. It can be found that the maximum adsorption capacity of SMS-10-5 for Pb²⁺ at 303.15 K are 50.5 mg g⁻¹ (Fig. 7a). A comparative study of our developed adsorbent (SMS-10-5) to other reported adsorbents toward Pb²⁺ removal was performed, and the results are presented in Table 1.^46–52 It can be seen that the adsorption capacity of SMS-10-5 toward Pb²⁺ ions is significantly higher than those found for many other conventional adsorbents, indicating that the SMS-10-5 has good adsorption performance for Pb²⁺.

Fig. 7 Isothermal adsorption data (a) and fitting curves (inset figures); kinetic adsorption data (b) and fitting curves (inset figure).

Furthermore, the analysis of the isotherm data by fitting them to different isotherm models was conducted to find the suitable model that could accurately depict these adsorption processes. Two common isotherm models (i.e., Langmuir and Freundlich models) were applied to fit the experimental data, respectively.⁵³ Clearly, the Langmuir model provides reasonably good fits to the experimental data (inset of Fig. 7a). Table 2 summarizes the obtained fitting parameters and correlation coefficients (R²). Results suggest that Pb²⁺ adsorption onto SMS-10-5 followed the Langmuir isotherm model. In other words, the surface of SMS-10-5 was typically homogeneous and Pb²⁺ was adsorbed onto the adsorbent in a monolayer manner.⁵⁴ Moreover, the values of equilibrium parameter R_L are found to be in the range of 0.058–0.429, indicating that the adsorption process of Pb²⁺ adsorption SMS-10-5 is considerably favorable.⁵⁵

Table 2 Adsorption isotherm constants for Pb²⁺ adsorption onto SMS-10-5

Langmuir isotherm model^a				Freundlich isotherm model^a
qm (mg g^−1)	KL (L mg^−1)	RL^b	R^2	KF (mg g^−1 (L mg^−1)^1/n)	1/n	R^2
a Ce is the equilibrium concentration of Pb^2+ (mg L^−1), qe is the amount of Pb^2+ adsorbed at equilibrium (mg g^−1), qm is the maximum adsorption capacity (mg g^−1), KL (L mg^−1) is the Langmuir binding constant, and KF (mg g^−1 (L mg^−1)^1/n) is the Freundlich constant.b , where C0 is the initial concentration of Pb^2+ (mg L^−1).
53.48	0.166	0.058–0.429	0.994	17.74	0.241	0.945

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In addition, the kinetic experiments of Pb²⁺ adsorption onto SMS-10-5 were also conducted, and the results are shown in Fig. 7b. The amount of Pb²⁺ adsorbed increases gradually by increasing the contact time from 2 to 240 min at 303.15 K. A large amount of Pb²⁺ ions is removed in the first 120 min, and the adsorption equilibrium is established in about 180 min. Compared with the other works, the equilibrium time is quite satisfied.⁵³

In order to further understand the adsorption process, two most widely used kinetic models were applied to fit the experimental data, which are pseudo-first-order and pseudo-second-order models, respectively.⁵³ Clearly, the pseudo-second-order model shows better fits to the experimental data than the pseudo-first-order one (inset of Fig. 7b). Table 3 summarizes the calculated rate constants and correlation coefficients (R²). The experimental q_e value (48.5 mg g⁻¹) agrees well with the calculated value (50.8 mg g⁻¹) using pseudo-second-order model, while the pseudo-first-order model fits poorly the experimental data. The obtained better correlation coefficient (R² = 0.992) by the straight-line plot of t/q_t against t also suggests that the adsorption of Pb²⁺ onto SMS-10-5 nanocomposite follows the pseudo-second-order model, indicating that the chemical adsorption is the rate-limiting step.⁵⁶

Table 3 Coefficients of pseudo-first-order and pseudo-second-order adsorption kinetic models

qe,exp^b (mg g^−1)	Pseudo-first-order model^a ln(qe − qt) = ln qe − k1t			Pseudo-second-order model^a
	k1 (min^−1)	qe,cal^c (mg g^−1)	R^2	k2 (g mg^−1 min^−1)	qe,cal^c (mg g^−1)	R^2
a qe and qt are the adsorption amount of Pb^2+ at equilibrium and at any time (mg g^−1), respectively; k1 (min^−1) and k2 (g (mg min)^−1) represent the equilibrium rate constants of pseudo-first-order and pseudo-second-order model, respectively.b Equilibrium adsorption capacity obtained from experiment.c Equilibrium adsorption capacity calculated according to kinetic models.
48.5	0.017	33.8	0.936	0.0011	50.8	0.992

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4. Conclusions

In the present work, we have demonstrated a practical route to the synthesis of waxberry-like SiO₂@MnSiO₃ core–shell nanocomposite. This process mainly involved a slow release of silicate species from silica sphere surface due to alkali etching and a surface precipitation between Mn²⁺ and silicate species, eventually leading to the formation of MnSiO₃ shell on the surface of silica nanosphere. Such a self-template method is facile, efficient, and possibly extendable to fabricate other core–shell nanocomposites. Taking advantages of the abundant and highly accessible active sites present in MnSiO₃ shell, the SiO₂@MnSiO₃ exhibited high activity in catalytic degradation of MB; more than 93% of MB could be decomposed within 40 min. Moreover, it was also proven to be a competent adsorbent for highly efficient removal of Pb²⁺ from wastewater; the Pb²⁺ adsorption capacity of SiO₂@MnSiO₃ was up to 50.5 mg g⁻¹.

Acknowledgements

The authors acknowledge the research grant provided by the Fundamental Research Funds for the Central Universities, China University of Geosciences (Wuhan) (No. CUG120115), Special Fund for Basic Scientific Research of Central Colleges, China University of Geosciences (Wuhan) (No. CUGL090305), Land Resources Geology Survey Projects of China (Grant No. 12120113015300), and National Natural Science Foundation of China (Grant No. 21303170).

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