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Article

Mn3O4 Catalysts for Advanced Oxidation of Phenolic Contaminants in Aqueous Solutions

1
Department of Chemical Engineering, Universitas Syiah Kuala, Banda Aceh 23111, Indonesia
2
ARC-PUIPT Nilam Aceh Universitas Syiah Kuala, Banda Aceh 23111, Indonesia
3
Department of Chemical Engineering, Universitas Riau, Pekanbaru 28293, Indonesia
*
Authors to whom correspondence should be addressed.
Water 2022, 14(13), 2124; https://doi.org/10.3390/w14132124
Submission received: 16 May 2022 / Revised: 25 June 2022 / Accepted: 28 June 2022 / Published: 3 July 2022
(This article belongs to the Special Issue Advances in Wastewater Resourcezation)

Abstract

:
Water-soluble organic pollutants, such as phenolic compounds, have been exposed to environments globally. They have a significant impact on groundwater and surface water quality. In this work, different Mn3O4 catalysts were prepared for metal oxide activation of peroxymonosulfate (PMS) to remove the phenolic compound from the water environment. The as-prepared catalysts were characterized using thermogravimetric-differential thermal analysis (TG-DTA), powder X-ray diffraction (XRD), scanning electron microscopy (SEM), and Brunauer-Emmett-Teller (BET) surface area analysis. Furthermore, the effect of temperature and reusability of the as-prepared Mn3O4 catalysts is also investigated. The Mn3O4 nanoparticles (NPs) catalyst reveals an excellent performance for activating PMS to remove phenol compounds. Mn3O4 NPs exhibits 96.057% efficiency in removing 25 ppm within 60 min. The kinetic analysis shows that Mn3O4 NPs fitted into pseudo-first order kinetic model and exhibited relatively low energy activation of 42.6 kJ/mol. The reusability test of Mn3O4 NPs displays exceptional stability with 84.29% efficiency after three-sequential cycles. The as-prepared Mn3O4 NPs is proven suitable for phenolic remediation in aqueous solutions.

1. Introduction

Rapid economic growth has benefited socio-economic development, but has increased the risk of environmental issues [1]. One of the most severe environmental problems is water pollution. The limited availability of freshwater, which only accounts for 2.5% of the planet’s water, faces water security threats that can affect 80% of the world’s population [2]. The emerging contaminants (ECs), such as organic pollutants, pesticides, and pharmaceuticals, have been increased through the mismanagement of industrial wastewater. Water-soluble organic pollutants, mainly phenolic compounds, are particularly widespread, with approximately 3 million tonnes identified in the environment globally [3]. The phenolic effluents come mostly from coal conversion processes, petroleum refineries, phenolic resin manufacturing, herbicide manufacturing, pharmaceuticals, pulp and paper, and petrochemicals [3,4]. Phenolic compounds are extremely harmful to humans due to their acute toxicity, mutagenicity, and bio-recalcitrant characteristic [3,5].
Different biological, chemical, and physical processes have been applied to remove organic pollutants from an aqueous solution, such as adsorption, biodegradation, coagulation-flocculation, ion exchange, photocatalytic, membrane electrodialysis and filtration, and solvent extraction. These strategies are incapable of completely removing phenolic compounds, require more extended time, and are not cost-effective [3,6,7,8,9,10,11]. Advanced oxidation processes (AOPs) have been proven to be a promising alternative to water treatment due to their compelling oxidation potential for the mineralization of organic pollutants [12]. This process generates radical species, such as the hydroxyl radical (OH) and the sulfate radical (SO4) [13]. Among those reactive radicals, the SO4 radical has higher reduction potential (E0 = 2.5–3.1 V vs. OH radical with E0 = 1.8–2.7 V), possesses a better half-life, better reaction stoichiometric efficiency than OH, and can be utilized in wider pH range [14,15].
SO4 radicals generally reproduce from the activation of peroxymonosulfate (PMS) or persulfate (PS) through heating, light irradiation, ultrasonication, or transition metals [14]. The activation through transition metals exhibits a distinguished advantage because it does not require external energy, unlike the other activations [16]. Various transition metals have been studied for phenolic compounds oxidation degradation, i.e., BiFeO3 magnetic nanoparticles [17], ZnO [18], Co-TiO2 [19], nano-Ag, Mn3O4-g-C3N4 [16], α-Mn2O3 [20], and dipicolinic acid-functionalized- Fe2O3 [21]. Of the various transition metals, Mn has superiority due to its abundance in nature, low toxicity, multivalence, inexpensiveness, and environmental friendly [16].
Moreover, the efficiency of Mn is influenced by different characteristics such as crystallinity, oxidation state and species, morphology, and specific surface area [16,22]. The effect of crystallinity demonstrated the most influence compared to other characteristics, e.g., MnO2 (146 m2/g) achieved an approximately eight times better performance compared to amorphous MnO2 (179 m2/g) [12].
In this present study, the evolution of morphology and degree of crystallinity of manganese oxides are investigated. The as-prepared transition metal oxide catalyst is used for sulfate radical generation using PMS, where the effect of morphology and degree of crystallinity for phenolic compound degradation is exhibited thoroughly. The investigation of the temperature effect and reusability is also presented using a kinetic analysis, where the estimated activation energy is a key parameter.

2. Materials and Methods

2.1. Materials

Potassium permanganate (CAS 7722-64-7), OXONE® (CAS 70693-62-8), polyethylene glycol (PEG200) (CAS 25322-68-3), and ethanol (CAS 64-17-5) were purchased from Sigma Aldrich, USA. All materials were used without any further purification.

2.2. Mn3O4 Catalysts Preparation

Mn3O4 950 °C was synthesized from as-prepared spherical MnO2 using the sonochemical method from Sankar et al., study [23]. Firstly, 0.8 g of potassium permanganate was dissolved in 50 mL of deionized water (DI water), then 10 mL of PEG200 was added to the aqueous solution. Afterward, the mixture was subjected to ultrasonication with a 20 kHz frequency for 15 min, where the reaction reached around 80 °C. The resulting brown precipitate was vacuum filtered using 0.45 μm Merck Millipore PVDF membrane and washed, firstly using DI water, and subsequently using ethanol. The obtained MnO2 was dried at 80 °C for 8 h. Finally, the as-prepared MnO2 was calcined at 950 °C. The as-prepared catalyst was denoted as Mn3O4 950 °C.
Secondly, the Mn3O4 NPs and octahedral catalysts were prepared using a hydrothermal method described by Li et al. [24], with some modifications. Firstly, Mn3O4 NPs was prepared using 0.105 g of potassium permanganate precursor, dispersed into 30 mL of DI water. Subsequently, 15 mL of PEG200 was added to the solution. The mixture was mixed for 30 min at ambient temperature before being put into a Teflon-line stainless steel autoclave. The mixture then underwent a hydrothermal process for 8 h at 120 °C. The as-prepared products were collected using vacuum filtration and washed several times using DI water and ethanol before drying at 60 °C for 12 h. The as-prepared catalyst was denoted as Mn3O4-NPs.
Lastly, octahedral Mn3O4 was prepared using a similar hydrothermal method with higher PEG200 concentrations. Five mol of potassium permanganate was dissolved into 30 mL of DI water with the addition of 30 mL of PEG200. Then, the as-prepared catalyst was denoted as octahedral Mn3O4.

2.3. Characterizations

The thermogravimetric analysis (TGA) (TGA/DSC1, Mettler Toledo, Columbus, OH, USA) was used to analyze the thermogravimetric profile of the MnO2. As-prepared catalysts were characterized for powder X-ray diffraction (XRD) using the Bruker-AXS D8 model (Bruker, Hannover, Germany) equipped with filtered Cu Kα radiation with a wavelength of 1.54178 Å and a scan range of 2θ of 5–70°. The surface morphology of the samples was examined using the NEON 40EsB model (ZEISS, Oberkochen, Germany). The average length and diameter of the samples were calculated using the image processing program ImageJ from 50 individual structures. Micromeritics Tristar 3000 (Micromeritics, Norcross, GA, USA) was used to acquire Brunauer–Emmett–Teller (BET) specific surface area, pore volume, and average pore radius.

2.4. Phenolic Degradation Performance

The typical catalytic oxidation of phenol was conducted in a 500 mL reactor containing 0.2 g/L as-prepared catalysts and immersed in 25 ppm, 250 mL phenol solutions. The reactor was dipped in a water bath, which had an attached temperature controller and stirring rod with a rotational speed of 400 rpm. Subsequently, after adsorption-desorption equilibrium is achieved, 2 g/L OXONE® was added to begin the oxidation process. At certain time intervals, 1 mL liquid solution was withdrawn from the reactor using a PVDF syringe filter of 0.45 μm. The liquid was then added to 0.5 mL methanol to prevent further reaction. The liquid concentrations were analyzed using high-performance liquid chromatography (HPLC), which is equipped with a UV detector at λ = 270 nm. The C-18 HPLC column was used with the mobile phase of 30% CH3CN and 70% ultrapure H2O with flowrate 1 mL/min. For chosen samples, total organic carbon (TOC) was defined applying a Shimadzu TOC-5000 CE analyzer.
The phenolic compound degradation efficiency was calculated using the following Equation (1):
Degradation   efficiency   ( % ) = C 0 C e C 0 × 100
where C0 and Ce is the initial and remaining phenol concentrations, respectively.
Furthermore, the kinetic analysis of the as-prepared samples were calculated using Langmuir-Hinshelwood (L-H) pseudo first-order kinetic equation as follow in Equation (2) [25]:
        ln ( C t C 0 ) = k × t
where C0 and Ct are the initial and final phenol concentration, respectively, k is rate constant (min−1), and t is time (min).
Additionally, the activation energy of the as-prepared catalysts was calculated at different temperatures (25–45 °C) using the Arrhenius equation. The Arrhenius equation, which correlated the relationship between the observed rate constant (kobs) and temperature, was presented in the following Equation (3) [26,27]:
ln   k o b s = ln   A E A RT
where A is the pre-exponential factor which can be calculated from the intercept of the Arrhenius plot, Ea. is activation energy (Kj.mol−1), T is temperature (K), and R is the gas constant (8.314 J/mol K).
Moreover, the reusability test was done with the solid catalyst from the previous experimental run. Before the next experimental run, the solid catalyst was washed thoroughly with DI water and dried at 60 °C for 12 h.

3. Results and Discussion

3.1. Thermogravimetric Profile of MnO2

The thermal stability and evolution profile of MnO2 are presented in Figure 1. The initial weight loss of 4.3361% in the range of 30–200 °C corresponds to the evaporation of adsorbed water content on the surface of MnO2 [28,29]. Moreover, observable two-phase weight losses were observed at 550 °C and 940 °C with 12.0075% and 16.2841% reduction, respectively. The first phase of weight loss at 550 °C is attributed to the evolution of MnO2 to Mn2O3 due to the oxygen release and reduction of Mn4+ ions into Mn3+ ions [30,31] The second phase of weight loss at 940 °C correlates with the sustained oxygen loss resulting in the conversion of Mn2O3 to Mn3O4 [30]. The TGA profile of Mn3O4 concurrence with previous studies that reported the thermal stability of MnO2 materials [30,32]. Furthermore, the transformation of as-prepared samples was examined further with XRD analysis. This TGA finding led to the choice of temperature for the hydrothermal synthesis of Mn3O4 catalysts at 950 °C.

3.2. Crystallinity Structure, Morphology, and Surface Area Analysis

The crystallinity of as-prepared Mn3O4 is observed in Figure 2a,b for Mn3O4 950 °C, Mn3O4 NPs, and octahedral Mn3O4. All 2θ diffractions of the as-prepared samples exhibit peaks at 18.13°, 29.03°, 31.14°, 32.41°, 36.21°, 38.12°, 44.56°, 50.94°, 54.06°, 56.18°, 58.64°, 60.01°, and 65.24° correspond to (101), (112), (200), (103), (211), (004), (220), (204), (105), (312), (224), and (400) planes, respectively. The standard diffraction is matched with JCPDS card No. 80–0382 correlated with the hausmannite spinel structure [24] The broad diffraction of the as-prepared samples possesses nanostructure dimensions, as seen in previous work [24]. Subsequently, Figure 2c exhibited different peak intensities of XRD spectra for the as-prepared samples. The Mn3O4 950 °C sample reveals the lowest intensity, followed by octahedral Mn3O4 and Mn3O4 NPs. The difference in peak intensity can be associated with different crystallite sizes and dislocation densities. The higher intensities correspond to larger crystallite sizes and better crystallinity. It is found that a larger crystallite size and higher crystallinity induced better degradation of organic compounds, as seen in a previous study [33]. Moreover, the determination of morphology of as-prepared Mn3O4 is presented using SEM imaging.
Figure 3a–c shows the SEM imaging of three as-prepared Mn3O4 samples. Figure 3a of Mn3O4 950 °C image shows clustered nanoparticles forming a quasi-spherical shape. The average diameter of the quasi-spherical shape of Mn3O4 950 °C is around 122.290 ± 43.0 nm. Moreover, the Mn3O4 NPs SEM image (Figure 3b) exhibits an agglomerated nanoplatelet with a length of approximately 85.211 ± 23.4 nm. These nanoplatelets clustered, forming pre-octahedral shapes. Finally, octahedral Mn3O4 SEM image (Figure 3c) reveals more dispersed nanoparticles with octahedral shapes, and slight nanorod shapes can be observed. The length of octahedral Mn3O4 is 0.65 ± 0.1 µm. Octahedron shapes could be formed utilizing accelerated reduction, and nucleation growth. However, the nanorod shapes could occur with excess PEG as a reducing agent and shape-directing agent [34].
Figure 4 illustrates the morphological evolution of Mn3O4 from nanoplatelet to octahedral architecture. Firstly, Mn3O4 forms a nanoplatelet-like shape. Afterward, the Mn3O4 is agglomerated, forming a pre-octahedral shape, depicted in Figure 3b highlighted with the red circles. Then, the Mn3O4 formed dispersed octahedral shapes (Figure 3c). The morphological evolution happened due to higher PEG200 concentrations. The polymer acts as a reducing agent and shape-directing agent [34]. The morphological evolution is followed by a self-assembly mechanism and then an Ostwald ripening mechanism [24].
The BET surface area, and N2 adsorption-desorption of each catalyst was analyzed to determine surface area, pore volume, and average pore radius. Table 1 lists the summary of the BET analysis. It reveals Mn3O4 950 °C possesses SBET 156.0 m2/g with a pore volume of 0.24 cm3/g and an average pore radius of 31.1 Å. The calculated surface area of Mn3O4 at 950 °C exhibited a higher BET surface area compared to other quasi-spherical Mn3O4 from previous studies, e.g., chemical leached Mn3O4 from manganese ore with a surface area of 9.32 m2/g [34], Mn3O4 synthesized using a precipitation method from potassium permanganate with a surface area of 21.2 m2/g [35], and solvothermal Mn3O4 from manganese(II)acetate tetrahydrate with a surface area of 131.49 m2/g [36]. Furthermore, different concentrations of the PEG200 alters the surface area of Mn3O4. Mn3O4 NPs exhibited higher surface area of 184.6 m2/g with an agglomerated platelet-like shape. With higher PEG200 concentrations, the Mn3O4 NPs transform from agglomerated nanoplatelets that resemble pre-octahedra into more polished octahedral shapes. Octahedral-shape controlled Mn3O4 was also exhibited in Li’s study, where the Mn3O4 formation evolution is exhibited through different reaction times [24]. The as-prepared octahedral Mn3O4 possesses a lower surface area compared to Mn3O4 NPs. This is due to the smaller size of individual nanoplatelets of Mn3O4 NPs compared to octahedral shapes of octahedral Mn3O4. This phenomenon can be seen from previous transformation morphology studies of Mn3O4 [24,34]. The calculated SBET of octahedral Mn3O4 is 122.4 m2/g is proven to be higher than octahedral Mn3O4 from previous works, i.e., hydrothermal Mn3O4 from KMnO4 with SBET of 23 m2/g [37] and hydrothermal Mn3O4 from manganese(II) acetate with SBET of 57.7 m2/g [38]. Moreover, all three as-prepared catalysts observed micropore-type pore size. This classification is determined by IUPAC classification with pore diameter of <2 nm, 2–50 nm, and > 50 nm classified as micropore, mesopore and macropore, respectively.

3.3. Phenolic Degradation Performance

The catalytic activity of the as-prepared Mn3O4 catalysts is evaluated from phenolic degradation using PMS presented in Figure 5. The initial PMS degradation without transition metal activation is performed to identify the contribution of PMS to phenolic degradation. It is found that PMS without activation degraded 2.465% phenolic compound, which has a negligible effect on the overall performance. This is because, with PMS alone, the sulfate radicals cannot be produced for phenolic oxidation [39]. Similar PMS without activation performance was also reported from previous work. Liu and Huang reported the role of PMS without activation is negligible, with only < 5% oxidation for the phenolic compound [40]. Furthermore, the preliminary performance of Mn3O4 without the presence of PMS is investigated using a Mn3O4 NPs sample. It is shown that the adsorption performance of Mn3O4 adsorbed 9.242% phenolic compound within 60 min. The low adsorption capacity of Mn3O4 is also recorded in previous studies. This occurs because phenol molecules easily penetrate through pores larger than 1 nm, while most NPs have pores smaller than this. Wang et al. reported no to little adsorption of phenol with only < 5% adsorption using Mn3O4 2D nanosheet after 2 h reaction time [5]. The higher adsorption rate we achieved in comparison to Wang et al. is due to the as-prepared Mn3O4 having a higher calculated surface area than in their work (65.2 m2/g).
The oxidation performance of PMS with different Mn3O4 catalyst activations is observed. All Mn3O4 catalysts are shown to efficiently activate PMS for phenol removal. Mn3O4 950 °C exhibited 49.730% efficiency within 60 min reaction time. While octahedral Mn3O4 reached 71.362% removal, the Mn3O4 NPs exhibited the best performance with 96.057% phenolic compound removal. This phenomenon occurred because the Mn3O4 NPs possess the most surface area and pore volume of the three catalysts, and these are the active sites where PMS activation occurs [12]. Additionally, octahedral Mn3O4 has the second-best oxidative degradation performance because it has a higher pore volume with a tighter pore radius, and therefore more active sites, compared to Mn3O4 950 °C. The greater availability of active sites for PMS activation is attributed to the fast-generating rate of sulfate radicals [39,41].
The two best-performing catalysts underwent investigation of different temperatures toward phenolic compound removal. Figure 6a shows octahedral Mn3O4 phenol removal performance at 25, 35, and 45 °C reaction temperatures. Moreover, the Langmuir-Hilselwood pseudo-first order kinetic model is fitted to determine the reaction rate [42]. Table 2 shows all the constant rates and the correlation coefficient of 0.99. The degradation result of octahedral Mn3O4 at 45, 35, and 25 °C all achieved 100% removal within 40, 120, and 190 min with kobs of 0.128, 0.073, and 0.030 min−1, respectively. Additionally, Figure 6b shows Mn3O4 NPs results at 45, 35, and 25 °C achieved 100%, 99.408%, and 96.057% removal within 40, 60, and 60 min with kobs of 0.186, 0.073, and 0.043 min−1, respectively. The result shows that higher temperature ensued better oxidation of phenol compounds. Higher temperature stimulates the chemical bond breakage and accelerates the decomposition of the persulfate, which leads to increased phenol removal [43,44,45]. TOC removal in Mn3O4 NPs/PMS systems was also measured and the results showed that roughly 78% of TOC removal was obtained for Mn3O4 NPs/PMS, at 190 min.
The correlation between temperature and reaction rate is further studied using the Arrhenius equation plot. Figure 7a,b show the Arrhenius plot for the as-prepared catalysts with a correlation coefficient of 0.99. It is found that octahedral Mn3O4 has an energy activation of 71.23 kJ/mol, which is much higher than Mn3O4 NPs with Ea of 42.6 kJ/mol. The energy activation of both catalysts is higher than diffusion-controlled reaction, which is in the range of 10–13 kJ/mol. This higher value indicates that the reaction rate is predominantly an intrinsic chemical reaction on the surface of oxide rather than a rate of mass transfer [46].
In addition, the comparison of PMS activation using several related catalysts for phenolic compound degradation is presented in Table 3. The as-prepared Mn3O4 NPs catalyst has relatively lower energy activation than several catalysts from previous works. Activation energy gives an insight into the minimum energy required for a chemical reaction to occur. The degradation of catalysts with lower activation energy starts faster compared to catalysts with higher activation energy [47].
Figure 8a,b depict the reusability test of octahedral Mn3O4 and Mn3O4 NPs evaluated in three sequential cycles. Both catalysts observed a decline in the efficiency trend. Mn3O4 NPs maintained an efficiency of 84.29% within 60 min after the third recycling, with only an 11.81% decline. On the other hand, octahedral Mn3O4 exhibited a 9.25% decrease in degradation performance with 87.35% efficiency within 150 min reaction time after the third recycle. The decline in efficiency can be attributed to the surface deactivation of the catalysts with intermediates on their surface [51]. However, the results of both catalysts show a considerably good performance after the third cycle since no additional purification is performed, except for simply washing them with water. Nevertheless, the best way to remove the intermediates from the catalyst surface is to calcinate and wash with ammonia solution, as Sun et al. reported [52].

3.4. Proposed Mechanism of PMS Activation by Mn3O4 Catalyst

The oxidation of organic compounds started when PMS contacted the active sites on the surface of Mn4. Then the active sites transferred a donor electron via a reduction-oxidation reaction, both Mn(III) and Mn(IV). Mn(IV) was reduced to Mn(III) and afterward oxidized to Mn(IV) using HSO5−. Subsequently, Mn(IV) activated PMS to produce SO4, then released into the bulk solution and interacted with H2O and OH to produce OH. The generated SO4 and OH would target phenol and produce numerous intermediates until the phenol is broken down into CO2 and H2O. The full step by step mechanism is illustrated in the following Equation (4) and Figure 9.
Equation (4) [5,41,53]:
Mn(4+) + HSO5 → Mn(3+) + SO4•− + OH
Mn(3+) + HSO5 → Mn(4+) + SO5•− + H+
SO4 + H2O → SO42 + OH + H+
SO4•− + OH¯ → SO42+ OH
OH + C6H5OH → Intermediates → CO2 + H2O + SO42
SO4•− + C6H5OH → Intermediates → CO2 + H2O + SO42

4. Conclusions

Three different Mn3O4 catalysts with different crystallinity, morphology, and specific surface areas were successfully synthesized using a simple hydrothermal method. The Mn3O4 NPs exhibit excellent performance in activating PMS for phenol removal with high efficiency and relatively low energy activation compared to other related catalysts. Mn3O4 NPs has an efficiency of 96.057% with Ea = 42.6 kJ/mol, followed by octahedral Mn3O4 with an efficiency of 71.362% with Ea = 71.23 kJ/mol, and Mn3O4 950 °C with the efficiency of 49.73%. The three-stage sequential recycling of Mn3O4 NPs shows a high degree of reusability, with only an 11.81% reduction in performance. This work provides a facile preparation of Mn3O4 NPs for the advanced oxidation process in water remediation applications.

Author Contributions

Conceptualization, S.M.; Data curation, S.M. and M.W.N.; Formal analysis, M.W.N.; Investigation, M.W.N.; Project administration, E.S.; Supervision, E.S.; Writing—original draft, S.M.; Writing—review & editing, E.S. and N.A. All authors have read and agreed to the published version of the manuscript.

Funding

This research received no external funding.

Institutional Review Board Statement

Not applicable.

Informed Consent Statement

Not applicable.

Data Availability Statement

Not applicable.

Acknowledgments

The authors would like to gratefully thank to the Chemical Engineering Department, Universitas Syiah Kuala and Chemical Engineering Department, Universitas Riau – Indonesia.

Conflicts of Interest

The authors declare no conflict of interest.

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Figure 1. TGA profile of MnO2.
Figure 1. TGA profile of MnO2.
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Figure 2. XRD pattern of (a) Mn3O4 950 °C (b) Mn3O4 NPs and octahedral Mn3O4, (c) combined XRD pattern of as-prepared Mn3O4 samples.
Figure 2. XRD pattern of (a) Mn3O4 950 °C (b) Mn3O4 NPs and octahedral Mn3O4, (c) combined XRD pattern of as-prepared Mn3O4 samples.
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Figure 3. Morphology of (a) Mn3O4 950 °C (b) Mn3O4 NPs (c) octahedral Mn3O4 catalysts.
Figure 3. Morphology of (a) Mn3O4 950 °C (b) Mn3O4 NPs (c) octahedral Mn3O4 catalysts.
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Figure 4. Morphological evolution of as-prepared Mn3O4 catalysts.
Figure 4. Morphological evolution of as-prepared Mn3O4 catalysts.
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Figure 5. Phenolic degradation efficiency in catalytic oxidation using manganese oxide catalysts. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L, and t = 25 °C.
Figure 5. Phenolic degradation efficiency in catalytic oxidation using manganese oxide catalysts. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L, and t = 25 °C.
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Figure 6. Effect of temperature on phenolic degradation in catalytic oxidation using (a) octahedral Mn3O4 (b) Mn3O4 NPs. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L.
Figure 6. Effect of temperature on phenolic degradation in catalytic oxidation using (a) octahedral Mn3O4 (b) Mn3O4 NPs. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L.
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Figure 7. Arrhenius plot of phenol degradation on (a) octahedral Mn3O4 (b) Mn3O4NPs catalyst.
Figure 7. Arrhenius plot of phenol degradation on (a) octahedral Mn3O4 (b) Mn3O4NPs catalyst.
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Figure 8. Phenol degradation efficiency of recycled (a) octahedral Mn3O4 (b) Mn3O4 NPs catalyst. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L.
Figure 8. Phenol degradation efficiency of recycled (a) octahedral Mn3O4 (b) Mn3O4 NPs catalyst. Reaction condition: phenol concentration = 25 ppm, catalyst = 0.4 g/L, PMS = 2 g/L.
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Figure 9. Schematic mechanism of PMS activation by Mn3O4 catalyst.
Figure 9. Schematic mechanism of PMS activation by Mn3O4 catalyst.
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Table 1. Surface area, pore volume and pore radius of a series of manganese oxides.
Table 1. Surface area, pore volume and pore radius of a series of manganese oxides.
CatalystSurface Area
(SBET, m2/g)
Total Pore Volume
(cm3/g)
Average Pore
Radius (Å)
Mn3O4 950 °C156.00.2431.1
Mn3O4 NPs184.60.778.01
Octahedral Mn3O4122.40.3318.16
Table 2. Kinetic constants of phenol degradation at different temperatures on α-Mn3O4 catalyst.
Table 2. Kinetic constants of phenol degradation at different temperatures on α-Mn3O4 catalyst.
CatalystTemperature (°C)kobs (min−1)R2Ea (kJ/mol)
Mn3O4 NPs250.0430.9842.60
350.0760.99
450.1860.99
Octahedral
Mn3O4
250.0300.9971.23
350.0730.99
450.1280.99
Table 3. Comparison of PMS activation using related catalysts for phenolic removal.
Table 3. Comparison of PMS activation using related catalysts for phenolic removal.
CatalystOperation ConditionEa (kJ/mol)Ref.
Mn3O4 nanorodsPhenol cons. 20 ppm, catalyst dosage 0.2 g/L, PMS dosage 6.5 mM59.7[5]
Cubic α-Mn2O3Phenol cons. 25 ppm, catalyst dosage 0.4 g/L, PMS dosage 2 g/L61.2[19]
MnO2/ZnFe2O4Phenol cons. 20 ppm, catalyst dosage 0.2 g/L, PMS dosage 2 g/L49.4[31]
Mesoporous biochar with modified KOH and CaCl2 (Ca/BS-800-KOH)Phenol cons. 20 ppm, catalyst dosage 0.066 g/L, PMS dosage 1 g/L68.3[43]
CoMnAl mixed metal oxide
(CoMnAl-MMO)
Phenol cons. 10 ppm, catalyst dosage 0.02 g/L, PMS dosage 1.5 g/L76.83[46]
CoMgAl layered double hydrotalcite (CoMgAl-LDH)Phenol cons. 0.1 mM, catalyst dosage 0.3 g/L, PMS dosage 0.3 mM65.19[48]
n3O4–reduced graphene oxide (Mn3O4-rGO)MO cons. 30 ppm, catalyst dosage 0.5 g/L, PMS dosage 1.5 g/L49.5[49]
Magnetic core/shell carbon nanosphere supported manganese (Mn/Air-MCS)Phenol cons. 20 ppm, catalyst dosage 0.2 g/L, PMS dosage 2 g/L59.5[50]
Octahedral Mn3O4Phenol cons. 25 ppm, catalyst dosage 0.4 g/L, PMS dosage 2 g/L71.23This work
Mn3O4 NPsPhenol cons. 25 ppm, catalyst dosage 0.4 g/L, PMS dosage 2 g/L42.60This work
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Muhammad, S.; Nugraha, M.W.; Saputra, E.; Arahman, N. Mn3O4 Catalysts for Advanced Oxidation of Phenolic Contaminants in Aqueous Solutions. Water 2022, 14, 2124. https://doi.org/10.3390/w14132124

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Muhammad S, Nugraha MW, Saputra E, Arahman N. Mn3O4 Catalysts for Advanced Oxidation of Phenolic Contaminants in Aqueous Solutions. Water. 2022; 14(13):2124. https://doi.org/10.3390/w14132124

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Muhammad, Syaifullah, Muhammad Wahyu Nugraha, Edy Saputra, and Nasrul Arahman. 2022. "Mn3O4 Catalysts for Advanced Oxidation of Phenolic Contaminants in Aqueous Solutions" Water 14, no. 13: 2124. https://doi.org/10.3390/w14132124

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