2.1 Materials

All reagents and solvents were purchased from commercial sources and were used without further purification.

2.2 Characterization

Transmission electron microscopy (TEM) was conducted on carbon-coated copper grids using a FEI Technai G2 F20 Field Emission scanning transmission electron microscope (STEM) at 200 kV (point-to-point resolution < 0.25 nm, line-to-line resolution < 0.10 nm). The TEM samples were prepared by placing 2–3 drops of dilute ethanol solutions of the nanomaterials onto carbon-coated copper grids. SEM was carried out with Philips CM120 and LEO 1430VP instruments. The pH of zero point charge, pH_ZPC, was determined by acid–base titration according to Davranche et al. [20]. The X-ray powder patterns were recorded with a Bruker D8 ADVANCE (Germany) diffractometer (Cu Kα radiation). Absorption spectra were recorded by a CARY 100 Bio VARIAN UV–vis spectrophotometer. FT–IR spectra were obtained by using a Unicam Matson 1000 FT-IR spectrophotometer using KBr disks at room temperature.

2.3 Synthesis of MnWO₄ nanoparticles

An aqueous solution (10 cm³) of Na₂WO₄·2H₂O (0.825 g; 2.5 mmol) was acidified with HCl solution (1 cm³, 6 M), and a white precipitate of WO₃·nH₂O (0.69 g) was obtained [21]. Then a solution of Mn(NO₃)₂·4H₂O (2.5 mmol) in 20 mL of deionized water was added to the obtained WO₃·nH₂O. This mixture was mechanically stirred and then H₂O was evaporated under vacuum using a rotary evaporator. After the third evaporation process, the left solid was ground and calcined at 300 °C for 6 h in the air.

2.4 Methylene blue degradation

The catalytic activity of the manganese tungstate nanoparticles was demonstrated by degrading MB in aqueous solution. First, a stock solution of aqueous MB (10 mg/L) was prepared. Fifty milliliters of aqueous MB were poured into a round-bottom flask, and 10 mg of catalyst were added to it. 0.1 mol of oxidant (TBHP) was added to the reaction mixture and was allowed to react in ambient conditions under stirring. For a given time interval, a small quantity of the mixture solution was pipetted into a quartz cell and its absorption spectrum was measured by a UV–visible spectrophotometer.

3.1 Catalyst characterization

In order to confirm the crystalline structure of the prepared MnWO₄, the XRD study was carried out (Fig. 1). According to the XRD pattern, pure crystalline MnWO₄ was obtained. All reflection peaks of the product can be easily indexed as a pure, monoclinic wolframite tungstate structure with space group P2/c, and the cell parameters of MnWO₄ were as follows: a = 0.48277 nm, b = 0.57610 nm, c = 0.49970 nm, α = γ = 90° and β = 91.14°, which is consistent with values taken from the literature (JCPDS Card Number: 80-0133) [22]. The sharpness of the diffraction peaks also indicates the high crystallinity of MnWO₄. Moreover, the main reflection peaks of all pure MnWO₄ have a trend to shift to low angles slightly compared to standard cards (on abscissa axis), which illustrates the increase in cell parameters.

To characterize the morphology of the prepared oxides, they were studied by scanning electron microscopy (SEM) and transmission electron microscopy (TEM). SEM and TEM images are shown in Fig. 2. The SEM images presented in Fig. 2a indicate the growth of sphere-like MnWO₄ crystals.

The TEM image of MnWO₄ shows approximately monodispersed nanocrystalline MnWO₄ particles with diameters of ca. 70–100 nm (Fig. 2b).

3.2 Catalytic effects

The decrease in percentage of dye concentration is a reliable criterion for the activity of the catalysts during the reaction, which is proportional to the disappearance of the dye color.

A plot of percentage decrease in [MB] versus time represents the influence of the oxidant on the decomposition of MB dye catalyzed by MnWO₄ nanoparticles at room temperature (Fig. 3). As it is shown, the removal of MB was not observed when only H₂O₂ or TBHP was added into the MB solution, which indicates that the direct oxidation of MB by H₂O₂ or TBHP was very limited and that no dye decomposition was observed, even after 120 min.

The effect of the TBHP amount on the percentage decomposition of MB was studied, and the results are shown in Fig. 4. When the amount of TBHP varies from 0 to 0.1 mol, the percentage decomposition of MB increased from 0 to ca. 90%.

When the MnWO₄ catalyst was added to a MB solution containing TBHP, the removal of MB was slightly increased, relative to the case when TBHP was alone in the system. However, as shown in Fig. 5. MB could be completely removed after 120 min in the presence of the catalyst and of TBHP, showing that MnWO₄ facilitates the oxidation of MB by TBHP. Although the use of WO₃ in the presence of TBHP can have a good influence on dye decomposition, our proposed method decomposes MB much more efficiently (Fig. 3).

It has been well established that the oxidative degradation of organic matter by metal oxides is accomplished via a surface mechanism [23], that is, the organic compound is adsorbed on the surface of metal oxides to form a surface precursor complex, then electron transfer occurs within the surface complex from the organic reductant to the surface-bound metal, followed by the release of organic oxidation products. The formation of a surface precursor complex is one of the kinetic rate-limiting steps in the heterogeneous oxidative degradation of organic pollutants, closely related to the nature of the surface charge.

The effect of pH on the degradation efficiency of MB was examined in the range from 3 to 9. As illustrated in Fig. 6, MB degradation worked very effectively over a wide pH range from 3 to 9, which is very favorable to the practical treatment of wastewater, since there is no need to adjust the solution pH. After 2 h, 51% of MB was degraded at pH 5 and an almost complete degradation occurred in neutral and alkaline solutions. The effect of pH on the catalytic reaction can be mainly explained by the surface charge of MnWO₄ (point of zero charge, pH_pzc of MnWO₄ ∼ 6.3). In aqueous solution, at pH values below 6.3, the catalyst's surface is positively charged and, above pH 6.3, it is negatively charged. Being a cationic dye, MB would prefer to adsorb on the negative surface. However, about 48% of MB can be degraded when the pH value of the solution approaches 3.

Fig. 7 shows the changes in the FT–IR spectrum of MB in the presence of a catalyst as a function of the reaction time. In curve a, the peaks at 1611 and 1407 cm⁻¹ are assigned to the CN and C–N bonds in the heterocycle of MB, respectively. As shown in curve b, these characteristic peaks have disappeared during the reaction time. Meanwhile, the appearance of new peaks at 1641, 1350 and 1210 cm⁻¹, attributed to the stretching vibrations of NO, N–O, and S = O bonds, can be seen.

It has been established that the complete oxidative degradation of methylene blue produces CO₂, HNO₃, H₂SO₄ and HCl [23]. In our system, H₂SO₄ and HCl were determined gravimetrically by precipitating them respectively as BaSO₄ and AgCl, using BaCl₂ and AgNO₃ as precipitants. As shown in Fig. 8, the SO₄²⁻ and Cl⁻concentrations increased with increasing the degradation time.

Finally, the effect of repeated uses of the catalyst on its catalytic activity was examined. After the end of the first experiment, the catalyst was washed with bi-distilled water, dried and a new experiment was started. The results in Fig. 9 demonstrate that about 90% of MB removal is accomplished after a five-cycle run in the presence of the MnWO₄ catalyst. It should be noted that the catalytic activity of the catalyst is almost unaffected upon repeated use.

We also found that the XRD patterns of both fresh and used catalysts are very similar (Fig. 10). The results show that the catalyst did not undergo any change during the reaction.