Heterogeneous Fenton-like degradation of methyl blue using MCM-41-Fe/Al supported Mn oxides

Abstract

Conventional heterogeneous catalysts often suffer from the constraints of high Fe³⁺ leaching concentrations and an amorphous appearance, which can affect the stability of the catalysts. We prepared uniform and spherical M-Fe/Al at room temperature. Furthermore, this material was supported with Mn oxides by an impregnation method to accelerate the degradation process of methyl blue. XRD, TEM, BET, and zeta potential measurements were employed to characterize the catalyst. In our work, M-Fe/Al–Mn showed a better methyl blue degradation efficiency than M-Fe/Al in the presence of H₂O₂. 93.3% discoloration was achieved by using M-Fe/Al–Mn within 200 min at 313 K. This catalyst performed better than M-Fe/Al. This could mainly be attributed to the increased adsorption effect of the Mn oxides. Furthermore, M-Fe/Al–Mn also showed good recyclability and stability over 5 cycles. Moreover, the Fe³⁺ leaching concentration was negligible (below 0.5 mg L⁻¹) in our work.

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1. Introduction

Recently, wastewater generated by the paper, food, cosmetic, and drug industries has accelerated environmental pollution. Unfortunately, toxic and structure-stable organic pollutants are still hard to treat with conventional methods.^1,2 AOPs (advanced oxidation progresses) have shown excellent catalytic activity for this kind of organic pollution. In particular, Fenton oxidation is powerful and extensively used in all AOPs.^3–5 It has attracted lots of attention since 1894. Furthermore, it has been widely accepted that the heterogeneous Fenton catalyst is more promising than the traditional homogeneous Fenton catalyst in practical applications because of its higher reusability and wider pH range tolerance.^6–9

Heterogeneous Fenton catalysts have been extensively researched by studying active metals and catalyst supports. Some metals (e.g. Fe, Cu, Ti, Mn, Co, and Ag^10,11) have been verified to have Fenton or Fenton-like catalytic activity. Among them, Fe has been investigated more, not only because of its inexpensive cost but also because of its high catalytic efficiency.^12–14 A survey of the literature indicated that heterogeneous Fe shows excellent catalytic activity in acidic systems. Besides, the carrier is another vital factor that affects the activity of a catalyst.^15–17 It is widely known that a large specific surface area and high hydrothermal stability ensure superior catalytic activity and stability. Much effort has focused on these aspects before.^18–20 MCM-41, a kind of mesoporous material, has been used as a carrier for heterogeneous Fenton-like catalysts due to its large specific surface area, open mesoporous structure and adjustable aperture.^21–24 Usually, the synthetic methods for catalysts based on MCM-41 include impregnation,^25,26 co-precipitation,^29,30 a chemical vapour deposition method,²⁷ etc. Among these, the co-precipitation process³⁵ is the main one and has attracted much attention due to the simple synthesis technique and high stability. Al-MCM-41, synthesized by a co-precipitation method, has been reported to exhibit an outstanding application prospect as a heterogeneous Fenton-like catalyst. The incorporation of Al(III) in the silica pore walls during the synthesis process achieves better uniformity of the pore size than without Al(III), which is beneficial for maintaining the stability of the catalyst.³¹ Moreover, Fe–, Cu– and Zn–Al-MCM-41 materials have been investigated too. This work showed that Cu could not be incorporated into the Al-MCM-41 framework. In contrast, Fe is sited in the framework of MCM-41.³² Fe–Al-MCM-41 showed a better stability and more active sites than Fe-MCM-41 showed. Incorporation of Al is advantageous for the recyclability and activity of the catalyst. Min Xia²² has used Fe–Al-MCM-41 as a Fenton-like catalyst for the degradation of phenol, and the possible beneficial role of Al has also been proposed. All in all, spherical Al-MCM-41 synthesized at room temperature by a co-precipitation method is a good choice of carrier.

There is no doubt that external power supplies, such as UV, ultrasound or microwaves, would enhance the Fenton-like process, which will block the practical application.²⁸ What’s more, increasing the adsorption capacity of carriers has been proven to be effective for accelerating the removal of organic pollution from solutions via the degradation process. In the study by Yimin Shao,³³ Mn/MCM-41 was confirmed to be a better adsorbent for the removal of methyl blue in aqueous solutions than MCM-41. Mn oxides on MCM-41 achieved an impressive promotion of the adsorption quantity of MCM-41. In this study, the promotion was mainly attributed to the electrostatic attraction effect of Mn oxides. Improving the adsorption capacity by the presence of Mn oxides on Al-MCM-41 has not been confirmed yet.

In our study, spherical MCM-41 was synthesized by introducing both iron and aluminum into MCM-41 by a co-precipitation method at room temperature.³⁴ Mn oxides were supported on MCM-41 using an impregnation method. As far as the authors know, spherical nanoparticles of MCM-41-Fe/Al supported with Mn oxides as a heterogeneous Fenton-like catalyst has been studied for the first time in this work. Subsequently, the degradation ability and reusability of M-Fe/Al–Mn and M-Fe/Al were also assessed in this study.

2. Experimental

2.1 Material

Hydrogen peroxide (H₂O₂, 30 wt%), methyl blue (MB), ferric nitrate (Fe(NO₃)₃·9H₂O), aluminum nitrate (Al(NO₃)₃·9H₂O), manganese acetate (Mn(AC)₂·4H₂O), tetraethylorthosilicate (TEOS), cetyltrimethylammonium bromide (CTAB), ammonia (25 wt%), etc. were of analytical grade unless otherwise stated.

2.2 Preparation of the catalyst

The samples were prepared by a co-precipitation method at 298 K using tetraethylorthosilicate (TEOS) as the Si precursor and cetyltrimethylammonium bromide (CTAB) as the template.^36,37 In a typical synthesis, CTAB (0.55 g), ferric nitrate (the molar ratio of Si and Fe was 100) and aluminum nitrate (the molar ratio of Si and Fe was 100) were dissolved in a mixed solution of 100 mL of deionized water and 35 mL of ethanol with stirring at 298 K for 5 min. Following this, 2.0 mL of TEOS was poured into the stirring solution until the solution was clear. Afterwards, 10 mL of ammonia (25%) was added to the system drop wise with continuous agitation to form a gel, which was stirred for 3 h. After that, the obtained gel was aged overnight. Finally, the sample was filtered, washed with deionized water, dried at 333 K in the oven for 24 h, and calcined at 823 K in a muffle for 5 h to completely remove the template from the sample. The obtained sample is expressed as M-Fe/Al. In this paper, M, M-Fe and M-Al represent pure MCM-41, Fe doped MCM-41 and Al doped MCM-41, respectively.

M-Fe/Al–Mn was prepared by an impregnation method using M-Fe/Al supported with Mn oxides (1.5 wt% Mn). More specifically, Mn(CH₃COO)₂·4H₂O was dissolved in 40 mL of deionized water. Then, 2.0 g of M-Fe/Al was added to the solution, which was stirred for 12 h. After that, the turbid liquid was dried at 353 K in the oven. Finally, the sample was calcined at 823 K in a muffle for 5 h to obtain M-Fe/Al–Mn. It is particularly worth mentioning that it was washed and dried before use.

2.3 Catalyst characterization

X-ray diffraction patterns of the samples were measured with a Rigaku MiniFlex 600, equipped with Cu Kα radiation, and operated at 40 kV and 200 mA. The shapes of the catalysts were mainly examined by using a TEM (JEM-2100). The nitrogen adsorption/desorption isotherms of the samples were collected with an analyzer (ASAP 2460) to learn about the surface area, pore diameter and pore volume. Furthermore, the zeta potentials were also measured by a Malvern potential analyzer to learn about the charge on the surfaces.

2.4 Reaction procedures

In our work, methyl blue was used as a model pollutant due to its highly stable structure and being easy to monitor. To study the different adsorption capacities of M-Fe/Al and M-Fe/Al–Mn, we put both the catalyst (0.5 g L⁻¹) and MB (20 mg L⁻¹) solution into a 50 mL conical flask, which was agitated using a shaking table, whilst maintaining a constant temperature of 298 K.

Degradation experiments were carried out in a 250 mL conical flask to assess the activity of the catalyst. Firstly, 250 mL of MB (10 mg L⁻¹) was added to the conical flask and the temperature was maintained at 313 K. Then, H₂O₂ solution was poured in, keeping the concentration between 0 and 20 mM. 0.1 M NaOH and HCl were used to adjust the pH value in the degradation process. Finally, 0.05 g of catalyst was added to the system to start the evaluation of the catalyst. Further research was conducted to study the stability of M-Fe/Al–Mn. The used catalyst was recycled by a centrifugal machine in our experiment. After that, the wet catalyst was transferred to an oven at 333 K for over 12 h. The dried catalyst was used for studying the recycling performance.

Adsorption experiments were conducted to compare the adsorption capacities of M, M-Fe/Al and M-Fe/Al–Mn on the shaking table. Both the MB solution and adsorbent were added to 20 mL test tubes at 298 K in a dark environment. After that, the test tubes were violently shaken on the shaking table. The concentration of MB was monitored by a UV-Vis spectrophotometer.

All the experiments were conducted in a dark environment to eliminate the influence of light. The Fe³⁺ leaching concentration was measured by a UV-Vis spectrophotometer. A UV-Vis spectrophotometer was also employed to monitor the concentration of MB (the maximum absorption wavelength of MB is 607 nm). The discoloration ratio (η) was calculated from the following eqn (1):

where c_t (mg L⁻¹) and c₀ (mg L⁻¹) represent the time-dependent concentration and the initial concentration, respectively.

3. Results and discussion

3.1 Catalyst characterization

X-ray diffraction is a powerful tool to characterize the structure of materials. The small and wide angle XRD patterns of the samples are presented in Fig. 1. In the small range (Fig. 1a), the XRD pattern of M exposed a typical MCM-41 structure with a strong diffraction peak at 2θ = 2.7°, which was indexed as a (1 0 0) diffraction. Unfortunately, the (1 1 0) and (2 0 0) diffractions of M were not strong, which could be attributed to the irregular inner structure. However, M-Al exhibited three strong diffraction peaks, which could be assigned as (1 0 0), (1 1 0) and (2 0 0) diffractions.

Fig. 1 (a) The low angle XRD patterns of different samples, and (b) the wide angle XRD patterns of different samples.

It should be noted that the (1 1 0) and (2 0 0) diffractions were even stronger than those of M at 2θ = 4.7° and 5.3°. There is no doubt that a more ordered structure could be gained by doping with Al. M-Fe showed a worse structure than M with only a weak diffraction peak, which could be indexed as a (1 0 0) diffraction. The (1 1 0) and (2 0 0) diffractions of M-Fe were unobservable. M-Fe/Al was more regular than both M and M-Fe. In other words, the XRD results suggested that the structure of M-Fe can be improved by doping with Al(III). M-Fe/Al was selected as a heterogeneous Fenton-like catalyst for our research. The wide angle XRD patterns (Fig. 1b) of M-Fe/Al and M-Fe/Al–Mn showed the change before and after loading of the Mn oxides. No obvious metal oxide diffraction peaks were observed, which indicated that Fe, Al and Mn might exist as monolayer oxides on the support or replace the position of Si in MCM-41, but they did not exist as a detectable metal oxide. Mn oxides would not change the structure of M-Fe/Al and would disperse well.

Both M-Fe/Al and M-Fe/Al–Mn exhibited well-ordered structures and spherical appearances by TEM (Fig. 2). Fig. 2a shows that M-Fe/Al was a three-dimensional spherical shape with particles 150 ± 50 nm in diameter, which is different from the traditional two-dimensional structure of MCM-41 and/or amorphous particles. The reason for this spherical morphology could be due to the use of ethanol, which is absent from the traditional synthetic process. The morphology of a growing particle depends on the rate of the polymerization of silicate and the rate of the mesostructure formation, which has been discussed by Chan et al. There is no doubt that the application of ethanol could minimize the surface free energy. Hence, a slower polymerization rate was obtained, which may break the traditional arrangement of micelles. Eventually, the shape of a sphere was obtained, as shown in the TEM picture. The TEM micrographs showed little difference between M-Fe/Al–Mn (Fig. 2b) and M-Fe/Al, which further confirms that the supporting Mn oxides hardly changed the shape or the structure of M-Fe/Al. This is consistent with the analysis of the wide angle XRD results. Remarkably, lattice fringes on the surface of MCM-41 are observable in Fig. 2c and d, which further verifies the existence of Mn oxides.

Fig. 2 TEM micrographs of the synthesized samples: (a) M-Fe/Al; (b–d) M-Fe/Al–Mn.

In order to figure out the valence states of the metals, M-Fe/Al–Mn was measured using the XPS technique, as shown in Fig. 3. Fe 2p_3/2 and Fe 2p_1/2 peaks were located at 712.4 eV and 726.7 eV, which indicated that Fe mainly existed as Fe(III). And so did Al, the characteristic peak of Al 2p was at 74.1 eV. For Mn, the main peaks for Mn 2p_1/2 and Mn 2p_3/2 were at 655.2 eV and 640.8 eV, respectively, which could be assigned to Mn(IV). In fact, Mn(IV) mainly existed as MnO₂ in M-Fe/Al–Mn.

Fig. 3 XPS spectra of M-Fe/Al–Mn: (a) for Fe 2p, (b) for Al 2p, and (c) for Mn 2p.

The nitrogen adsorption/desorption isotherms of M-Fe/Al–Mn and M-Fe/Al are illustrated in Fig. 4. Both of the samples exhibited type IV N₂ isotherms, which is characteristic of a typical mesoporous material. This result indicated that after loading the Mn oxides, the structure of MCM-41 almost remained intact. The detailed data on the BET surface areas of M-Fe/Al and M-Fe/Al–Mn are presented in Table 1. According to the results, the BET surface area of M-Fe/Al–Mn was a little smaller than that of M-Fe/Al. The pore diameter and pore volume showed a similar trend, which could be attributed to the existence of Mn oxides. M-Fe/Al–Mn showed a large specific surface area, which could guarantee degradation efficiency.

Fig. 4 Nitrogen adsorption isotherms of M-Fe/Al and M-Fe/Al–Mn.

Table 1 The surface areas, pore diameters and pore volumes of M-Fe/Al and M-Fe/Al–Mn

Sample	Surface area (m^2 g^−1)	Pore diameter (nm)	Pore volume (cm^3 g^−1)
M-Fe/Al	1410.26	2.97	2.52
M-Fe/Al–Mn	1105.96	2.44	1.99

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3.2 Adsorption of MB

In order to investigate the adsorption capacity of M-Fe/Al, M-Fe/Al–Mn and M, we carried out adsorption experiments under the same experimental conditions, ensuring the ratio of MB/catalyst was 1/10. The results are shown in Fig. 5. It is obvious that the adsorption process was nearly established within 50 minutes. Then, the rate of adsorption slowed down gradually. The adsorption and desorption process achieved a balance after 200 minutes. M-Fe/Al–Mn achieved a better MB adsorption efficiency than M or M-Fe/Al. The final adsorption quantity of M-Fe/Al–Mn was about 45 mg g⁻¹, which was much lager than that of M-Fe/Al or M. The increase in the adsorption quantity and adsorption rate is mainly attributed to the effect of Mn oxides, which might be explained by the effect of electrostatic attraction. In other words, the incorporation of Al in M-Fe–Mn does not affect the electrostatic attraction for the promotion of adsorption. Moreover, the phenomenon of electrostatic attraction was confirmed by the zeta potentials. The zeta potentials of M-Fe/Al–Mn and M-Fe/Al were −14.4 mV and −17 mV, respectively, at pH = 7. However, the surface of the Mn oxide particles was positively charged, which will make M-Fe/Al–Mn less negative than M-Fe/Al, which goes against the adsorption of MB. In conclusion, the presence of Mn oxides on Al-MCM-41 could improve the adsorption capacity efficiently.

Fig. 5 The adsorption performance of M-Al, M-Fe/Al and M-Fe/Al–Mn (adsorption dosage = 0.2 g L⁻¹, MB = 20 mg L⁻¹, pH = 7, temperature = 298 K).

3.3 Degradation of MB

3.3.1 The effect of the initial pH value. The effect of the initial pH value plays an important role in the degradation process. We studied the removal of MB by using M-Fe/Al–Mn with different initial pH values ranging from 3 to 8. The results are presented in Fig. 6. In the first 40 minutes, the absolute value of the slopes was large, which may be related to the adsorption of the catalysts. In addition, M-Fe/Al–Mn obtains a good degradation rate in acidic solution systems. The decolorization rate of MB mainly decreased when the pH value increased. It was obvious that an acidic environment would be in favour of the Fe redox process, which could accelerate the degradation process. It is noteworthy that the degradation results were similar when the pH value was 4 or 3. The concentration of leached Fe³⁺ would be higher when the pH value is 3 compared to when the pH value is 4. A pH value of 3 is not conducive to maintaining the stability of the catalyst. Therefore, a pH value of 4 was chosen as the optimum pH value for our further experiments.

Fig. 6 The effect of the initial pH value on the decolorization of MB (catalyst dosage = 0.2 g L⁻¹, MB = 5 mg L⁻¹, H₂O₂ = 0.1 g L⁻¹, temperature = 333 K).

3.3.2 The effect of the concentration of H₂O₂. The concentration of H₂O₂ is directly related to the generation of the hydroxyl radical, which can oxidize organic pollutants. Hence, we conducted experiments using M-Fe/Al–Mn as the catalyst to study the effect of the initial hydrogen peroxide concentration and the results are shown in Fig. 7. The degradation efficiency of MB increased with the increase in H₂O₂ concentration. However, when the concentration of H₂O₂ was as high as 20 mM, the degradation efficiency decreased unexpectedly. This is because a higher H₂O₂ concentration could scavenge more hydroxyl radicals, resulting in a low degradation efficiency. In our study, 10 mM H₂O₂ is appropriate for the degradation experiments.

Fig. 7 The effect of the initial H₂O₂ concentration on the degradation of MB (catalyst dosage = 0.2 g L⁻¹, MB = 5 mg L⁻¹, temperature = 333 K, pH = 4).

3.3.3 Catalytic activity. The discoloration results are shown in Fig. 8. H₂O₂ showed no catalytic ability towards MB when used alone. Due to the adsorption properties of M-Fe/Al–Mn, the removal of MB was only 10%. But with H₂O₂, the discoloration of MB with M-Fe/Al–Mn could reach 93.3%. With the same concentration of H₂O₂, the degradation result of M-Fe/Al was only 80.0%, which showed a lower degradation ability than M-Fe/Al–Mn. Moreover, the COD (Chemical Oxygen Demand) removal result in the presence of M-Fe/Al–Mn and H₂O₂ was also studied to make sure that MB was indeed degraded. After 200 min, the COD removal result obtained 30%. This confirmed that part of MB had been oxidized. The possible effect of leached Fe³⁺ (0.4 mg L⁻¹) was also taken into consideration – the decolorization rate of MB was 15% after 200 min by using 0.4 mg L⁻¹ Fe³⁺ under the same experimental conditions.

Fig. 8 Discoloration and COD removal of MB under different experiment conditions (catalyst dosage = 0.2 g L⁻¹, MB = 5 mg L⁻¹, H₂O₂ = 0.1 g L⁻¹, temperature = 313 K, pH = 4).

The UV-Vis absorption changes over time were measured (Fig. 9) to study the structural and molecular features of MB during the degradation process by using M-Fe/Al–Mn as the catalyst. It was obvious that there were two characteristic bands. One main band in the visible region was observed at 605 nm, which originates from the extended chromophore. The absorption intensity decreased very fast, which is consistent with the results in Fig. 8. Another characteristic band was at 304 nm in the UV region, which could be assigned to benzene-like structures in the molecule. However, the benzene-like structures were more stable than the extended chromophore. Hence, the degradation rate of phenyl was slower. After 30 min, the absorption of MB was negligible, which indicated that the benzene-like structures and extended chromophore were degraded to smaller aliphatic compounds and eventually CO₂ and H₂O.

Fig. 9 UV-vis spectra of a typical degradation process (catalyst dosage = 0.2 g L⁻¹, MB = 10 mg L⁻¹, H₂O₂ = 0.1 g L⁻¹, temperature = 353 K, pH = 4).

3.4 Stability of M-Fe/Al–Mn

The reusability performance of M-Fe/Al–Mn is shown in Fig. 10a. The discoloration rate of MB reached nearly 90% within 50 minutes after five cycles. Moreover, M-Fe/Al–Mn exhibited perfect recyclability. Furthermore, leaching tests were also conducted to evaluate the concentration of Fe³⁺ in the solution after every degradation experiment at 333 K. In Fig. 10b, the leaching experiments showed that the leached Fe³⁺ concentrations were both below 0.5 mg L⁻¹. Remarkable, the leached Fe³⁺ concentration of M-Fe/Al–Mn was lower than that of M-Fe/Al. The leaching experiments provided evidence to verify the excellent stability and recyclability of M-Fe/Al–Mn, which could ensure the actual application value.

Fig. 10 (a) Recycling experiments for M-Fe/Al–Mn (catalyst dosage = 0.2 g L⁻¹, MB = 5 mg L⁻¹, H₂O₂ = 0.1 g L⁻¹, temperature = 333 K, pH = 4), and (b) the leached Fe³⁺ concentration after every experiment (catalyst dosage = 0.2 g L⁻¹, MB = 5 mg L⁻¹, H₂O₂ = 0.1 g L⁻¹, temperature = 333 K, pH = 4).

3.5 Reaction mechanism

The degradation process of MB is shown below based on the heterogeneous Fenton reaction, which is widely recognized.

≡Fe²⁺ + H₂O₂ → Fe²⁺ + HO˙ + OH⁻ (2)

≡Fe³⁺ + H₂O₂ → ≡Fe²⁺ + ˙O₂H + H⁺ (3)

MB + ˙OH → degradation products (CO₂ and H₂O) (4)

The possible degradation process of MB in our work based on iron redox cycling is proposed in Fig. 11: the negative ion of MB is adsorbed onto the surface of MCM-41 before degradation. Fe(III) and Al(III) replace some positions of Si(IV), which results in more negative charge on the catalysts to inhibit the adsorption of the MB negative ions. On the other hand, Mn oxides show a positive charge in aqueous solutions. In summary, M-Fe/Al supported with Mn oxides would make the catalyst exhibit a relatively positive charge. It is reasonable that using M-Fe/Al–Mn as an adsorbent could obtained better adsorption of MB than M-Fe/Al or M. The results from the MB adsorption curve and the zeta potential data verify the theory. Meanwhile, M-Fe/Al–Mn presented the best degradation result, as shown in Fig. 8. The promotion of degradation may be attributed to the increased adsorption effect of the Mn oxides on the surface of M-Fe/Al. Firstly, the hydroxyl radical is produced by heterogeneous Fe(III), whose concentration was higher on the surface of M-Fe/Al–Mn than in the solution. The concentration of the MB ions on the surface of M-Fe/Al–Mn is also larger than that in the solution because of electrostatic attraction, which increases the chance of them reacting with hydroxyl radicals and results in a better degradation performance of MB.

Fig. 11 Proposed mechanism for the degradation acceleration of MB using M-Fe/Al–Mn.

4. Conclusions

In our work, we combine both co-precipitation and impregnation methods to synthesize M-Fe/Al–Mn. The spherical M-Fe/Al catalyst was synthesized at room temperature by a co-precipitation method. Mn oxides were supported on M-Fe/Al by an impregnation method. M-Fe/Al–Mn showed a typical three-dimensional spherical morphology with particles 150 ± 50 nm in diameter. Even the specific surface area of M-Fe/Al was larger than that of M-Fe/Al–Mn. M-Fe/Al–Mn showed better MB degradation results, which could be attributed to the electrostatic attraction of Mn oxides. Moreover, the Fe³⁺ leaching concentration was less than 0.5 mg L⁻¹ and the catalyst behaved almost the same over 5 cycles. M-Fe/Al–Mn has been proven to be an attractive catalyst and the application of Mn oxides reveals a new strategy for enhancing the activity of heterogeneous Fenton-like catalysts.

Acknowledgements

The authors are grateful to the National Major Science and Technology Program for Water Pollution Control and Treatment (No. 2013ZX07210001) and Engineering technology research of multiple-effect desalination in low temperature (DZY-150001-KY-F3).

Notes and references

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