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Heterogeneous degradation of organic pollutants by persulfate activated by CuO-Fe3O4: mechanism, stability, effects of pH and bicarbonate ions YANG LEI, Chuh-Shun Chen, Yao-Jen Tu, Yao-Hui Huang, and Hui Zhang Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b00623 • Publication Date (Web): 08 May 2015 Downloaded from http://pubs.acs.org on May 9, 2015
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Environmental Science & Technology
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Heterogeneous degradation of organic pollutants by
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persulfate activated by CuO-Fe3O4: mechanism, stability,
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effects of pH and bicarbonate ions
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Yang Lei a, b, Chuh-Shun Chen b, Yao-Jen Tu c, Yao-Hui Huang b, d*, Hui Zhang a*
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a
Department of Environmental Engineering, Wuhan University, Wuhan.
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b
Department of Chemical Engineering, National Cheng Kung University, Tainan.
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c
Institute of Urban Study, Shanghai Normal University, No.100 Guilin Rd. Shanghai.
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d
Sustainable Environment Research Center, National Cheng Kung University, Tainan.
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*Corresponding author
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Phone: 86-27-68775837;
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Fax:
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E-mail:
[email protected]. (H. Zhang)
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Phone: 886-2757575x62636;
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Fax:
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E-mail:
[email protected]. (Y.H. Huang)
86-27-68778893;
886-6-2344496;
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Graphical abstract
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ABSTRACT
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Magnetic CuO-Fe3O4 composite was fabricated by a simple hydrothermal method and
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characterized as a heterogeneous catalyst for phenol degradation. The effects of pH
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and bicarbonate ions on catalytic activity were extensively evaluated in view of the
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practical applications. The results indicated that an increase of solution pH and the
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presence of bicarbonate ions were beneficial for the removal of phenol in the
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CuO-Fe3O4 coupled with persulfate (PS) process. Almost 100% mineralization of 0.1
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mM phenol can be achieved in 120 min by using 0.3 g/L CuO-Fe3O4 and 5.0 mM PS
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at pH 11.0 or in the presence of 3.0 mM bicarbonate. The positive effect of
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bicarbonate ion is probably due to the suppression of copper leaching as well as the
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formation of Cu(III). The reuse of catalyst at pH0 11.0 and 5.6 showed that the
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catalyst remains a high level of stability at alkaline condition (e.g. pH0 11.0). On the
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basis of the characterization of catalyst, the results of metal leaching and EPR studies,
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it is suggested that phenol is mainly destroyed by the surface-adsorbed radicals and
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Cu(III) resulting from the reaction between PS and Cu(II) on the catalyst. Taking into
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account the widespread presence of bicarbonate ions in waste streams, the
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CuO-Fe3O4/PS system may provide some new insights for contaminant removal from
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wastewater.
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INTRODUCTION
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With the decrease of freshwater resources, recovery of water from wastewater is
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increasingly important for the sustainable development of the world.1 Among the
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various water remediation technologies, advanced oxidation processes (AOPs), which
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are based on the generation of reactive oxygen species (ROS), are regarded as an
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effective technology for the degradation of hazardous organic pollutants in
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wastewater.2, 3 Therefore, much effort has been invested to improve the treatment
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efficiency of AOPs as well as extending their working conditions and cutting down on
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secondary pollution. In contrast to homogenous AOPs, heterogeneous AOPs have
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received extensive attention as they can usually work in mild conditions and produce
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almost no by-product sludge after reaction.4, 5
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Catalysts have played a very important role in heterogeneous AOPs with iron
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oxides (e.g., Fe3O4 and γ-Fe2O3)6, 7 and iron-immobilized silica4, 8 being particularly
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studied due to the fact that they are highly accessible because iron is the second most
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abundant element in the earth’s crust.9 For example, our group previously fabricated
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shape-controlled nano Fe3O4 as an efficient Fenton-like catalyst for the mineralization
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of phenol.7 However, many of those catalysts are often confronted with weak catalytic
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activity, which often needs the aid of electricity, ultrasonic sound and UV or visible
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light irradiation.10, 11 However, a drawback is that the introduction of external energy
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increases the cost of the treatment process. Alternatively, an efficient way to improve
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the catalytic activity is to replace some activate sites of the iron oxides with other
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metal ions. Accordingly, the new formed catalysts, named (Co, Cu, Zn)Fe2O4, have
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been the focus of many studies in past decades.12-18 The bimetallic catalyst, CuFe2O4
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has found the most favor, not only because of its good catalytic performance but also
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copper is not regarded as a potential carcinogen.19
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Some researchers have reported that the CuFe2O4 activated peroxymonosulfate
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(PMS) process is quite effective for the destruction of organic contaminants in
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water.15-17 However, it was also concluded by Guan et al.17 that only PMS could be
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easily activated by CuFe2O4. Indeed, the asymmetric structure of PMS appeared to
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makes it more readily activated than persulfate (PS) and hydrogen peroxide (H2O2)
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which have symmetrical structures.20 However, as is well-known, PS is the most
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suitable of the three common used oxidants in AOPs because it is cheaper than PMS
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and more stable than H2O2 while have similar oxidation ability after activation.20, 21 In
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order to effectively activate PS, it is therefore necessary to develop new catalysts.
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Recently, the activation of PS by copper oxide particles was reported whereby
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Liang et al.22 explored the degradation of p-chloroaniline by copper oxidate activated
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PS and concluded that different radicals were involved at various solution pHs.
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However, Zhang et al.23 reported that the CuO/PS coupled system could selectively
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attack some organic pollutants but did not rely on the generation of radicals. Thus,
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apparently, there is some controversy about the mechanism of copper oxide activated
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PS process. On the other hand, although copper oxide showed good performance
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under laboratory conditions, the leaching of hazardous copper ions may produce
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separate problems in industrial applications. From the viewpoint of the possible
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application of copper based AOPs for wastewater treatment, it is therefore urgent to
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solve these problems, as well as determining the mechanism of the CuO activated PS
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process.
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A possible way to overcome these drawbacks is to combine CuO with Fe3O4, which
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simultaneously utilizes the high catalytic property of CuO and the magnetic property
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of Fe3O4. Therefore, in this current study, CuO-Fe3O4 was fabricated using a simple
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one-step method and its role as a heterogeneous catalyst for activating PS was studied.
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Specially, the effect of pH on the coupled process (CuO-Fe3O4/PS) was especially
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studied. The primary reason why solution pH greatly affected the performance and the
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stability of CuO-Fe3O4 may be mainly due to the leaching of copper ions, rather than
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the electrostatic interaction between the substrate and catalyst, which had been
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suggested previously.16,
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water and wastewater, the influence of NaHCO3 in CuO-Fe3O4/PS system was also
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studied. Interestingly, it was found that the coupled system performed well at higher
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pH and even better in the presence of bicarbonate. Finally, based on the
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characterization of the catalyst and the experimental results, the mechanism of
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CuO-Fe3O4 activated PS process is clearly elucidated.
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MATERIALS AND METHODS
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In addition, as bicarbonate is widely present in natural
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Chemicals. Phenol and methanol were obtained from Sigma-Aldrich, Taiwan.
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Sodium persulfate and sodium bicarbonate were bought from Merck, Taiwan. The
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5,5-dimethyl-1-pyrrolidine N-oxide (DMPO) was purchased from Aladdin, China. All
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reagents used were at least analytical grade and prepared in Milli-Q water.
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Preparation and Characterization of Catalyst. The CuO-Fe3O4 catalyst was
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fabricated using a simple and one-step hydrothermal method. CuSO4·5H2O (500 mL,
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0.1 M) and FeSO4·7H2O (500 mL, 0.5 M) were mixed under continuous air purging
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(flow rate = 3 L/min). During the reaction, the pH and temperature were kept constant
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at 8.0 and 80 °C, respectively. The fabrication process can be simply expressed as
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follows:
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3Fe2+ + 6HO− + 0.5O2 → Fe3O4 + 3H2O
(1)
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Cu2+ + 2HO− → CuO +H2O
(2)
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The desired solids were separated with a magnet and then washed 3 times with 1.5
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L distilled water and dried at 105 °C. Various techniques such as Fourier transform
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infrared (FTIR, Bruker Tensor 27), X-ray powder diffraction (XRD, Rigaku, RX III),
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scanning electron microscopy coupled to energy dispersive spectrometer (SEM-EDS,
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JEOL JSM-6700F), and Brunauer–Emmett–Teller (BET) sorptometer were used to
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characterize the catalyst. X-ray photoelectron spectroscopy (XPS, ESCALAB 250Xi)
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was conducted for the determination of the chemical species of copper and iron.
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Experimental Procedure. Batch trials were performed in glass bottles with phenol
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solution (400 mL, 0.1 mM). The catalyst (0.12 g) was added, followed by the PS (5.0
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mM) to start the reaction. Where required, the initial pH (pH0) was adjusted by
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addition of appropriate amounts of H2SO4 (0.1 M) or NaOH (0.1 M). During the
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investigation of the effect of NaHCO3 appropriate amounts of NaHCO3 were added
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before addition of PS. Samples were removed at predetermined time intervals and
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were filtered through 0.22 µm syringe filters prior to analysis. Unless specifically
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noted, all the experiments were carried out at an ambient temperature of 20 ± 2 °C
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and were exposed to air.
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EPR Studies. For Electron Paramagnetic Resonance (EPR) studies, experiments
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were conducted using DMPO as a spin-trapping agent. A solution containing 10 mM
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DMPO, 5.0 mM PS was prepared at pH = 5.6±0.1, and then catalyst was added to
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initiated the reaction. After 5 min of reaction, samples were taken and analyzed on a
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JEOL FA200 spectrometer at room temperature. EPR measurements were conducted
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using a radiation of 9.147 GHz (X band) with a modulation frequency of 100 kHz,
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modulation width of 0.1 mT, a sweep width of 20 mT, center field of 326.0 mT, scan
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time of 60 s, time constant of 0.03 s, and microwave power of 5 mW.
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Analysis. The phenol concentration were determined with high performance liquid
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chromatograph (HPLC, Shimadzu 6A) equipped with a TSK-GEL ODS-100S column
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(4.6 mm × 250 mm) and a UV detector at wavelength of 270 nm. An eluent of
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methanol and water (V/V, 70/30) was used as the mobile phase with a constant flow
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rate (0.8 mL/min). The injection volume was 20µL and the retention time was 5.2 min.
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The mineralization of phenol was measured with a TOC Analyzer (Sievers 900, GE).
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The zero point charge (pHzpc) of CuO-Fe3O4 was determined by mass titration.17 The
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metal ions were quantified by inductively coupled plasma optical emission
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spectrometer (ICP-OES, ULTIMA 2000, HORIDA).
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RESULTS AND DISCUSSION
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Characterization of the CuO-Fe3O4 catalyst. Figure 1(A) shows the XRD pattern
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of CuO-Fe3O4 at 2θ from 10-70o. The sharpness of XRD reflections clearly illustrates
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that the synthesized catalyst is highly crystalline. The well matched peaks with
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standard Fe3O4 (JCPDS: 65-3107) and CuO (JCPDS: 45-0937) demonstrate that the
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fabricated catalyst is CuO-Fe3O4. The EDS results [see Figure 1(B)] show that the
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catalyst is mainly formed of the three elements, O, Cu and Fe. In addition, the
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atomic% ratio of Fe to Cu is 4.7, which is very close to the designed value (5.0).
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Furthermore, the SEM photo [Figure 1(B)] shows that the size of the fabricated
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CuO-Fe3O4 was at nano level and this was further confirmed by BET analysis (Table
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S1). These results clearly show that nano sized CuO-Fe3O4 was successfully prepared
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by the simple one-step method.
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Phenol Degradation. As can be seen from Figure 2, while phenol was not removed
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by PS or CuO-Fe3O4 alone, phenol was effectively degraded when they were used
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together. The removal and mineralization percentages of phenol were 80% (Figure 2)
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and 78% (Figure S1) in 120 min, respectively. All these evidences suggest that
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CuO-Fe3O4 coupled with PS is an effective system for the degradation of organic
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pollutants from water.
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Concentrations of leached metals were also monitored in this process. Considering
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the solution at pH = 5.6±0.1, Figure 2 shows that there was almost no iron leached
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while the amount of copper ions in solution gradually increased over the 120 min of
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the experiment. To check the effect of leached Cu2+, control experiments with Cu2+
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activated PS were performed. Figure 3 indicates that no phenol was degraded at the
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leached Cu2+ concentration (5.6 mg/L), and no phenol was removed even when the
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dose of Cu2+ was increased to 64 mg/L. These results showed that the potential effect
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of leached ions was negligible and also demonstrated that the degradation process
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only occurred at the surface of catalyst. Moreover, the heterogeneous reaction
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mechanism was supported by a 44% inhibition of phenol removal when a magnet was
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used to attract CuO-Fe3O4 (Figure S2).
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Effect of pH. It is well known that the performance of homogeneous AOPs is
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strongly correlated to the solution pH.24 Recently, some studies that focused on copper
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oxide activated PS and CuFe2O4/PMS systems also found that the processes were
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influenced by pH.16, 17, 22 Therefore, the degradation of phenol at different solution pH
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was explored and Figure 4(A) shows that the degradation of phenol is obviously pH
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dependent. There was almost no phenol removed at pH0 = 2.5. The amount removed
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increased dramatically, reaching 68%, when pH0 = 4.0 and then smoothly increased
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with the increase of pH0 before it reached 8.5 (80%). Thereafter, there was a second
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apparent jump when the pH0 was increased to 10.0. The phenol removed was further
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increased if the pH0 = 12.0 (98%). On the other hand, when the reaction rate constant
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in the first stage (10 min) of removal was calculated (Figure S3), it was found that the
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reaction rate constant increased by 5 times when the pH0 was increased from 4.0 to
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12.0, as seen by rates of 0.001 and 0.005 mM/min, respectively. Thus, the effect of
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pH0 on the degradation of phenol is clear, the higher the solution pH, the higher the
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removal rate. Nevertheless, as stated, the electrical interaction between catalyst and
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pollutants may not fully explain the results in our system. For instance, the electrical
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interaction between CuO-Fe3O4 and phenol (pKa = 10) should be ignored at the pH0
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range (2.5-7.0), as the pHzpc of the CuO-Fe3O4 is about 7.3. Furthermore, at pH0
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above 10, the phenol is present in its deprotonated form and the catalyst is also
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negatively charged, which means electrical repellent is dominant between phenol and
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CuO-Fe3O4. Thus, it is speculated that there must be other factors rather than
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electrical interaction that largely affect the degradation process. In order to better
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understand the impact of pH, the change of solution pH with time elapse was
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monitored and the concentration of leached free metal ions was measured for all pHs
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studied.
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Interestingly, it was found that the removal trend of phenol at various pH0
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corresponds to the change of solution pH during the reaction, as well as the leaching
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of Cu2+. Figure 4(B) shows that the pH values remain unchanged after 40 min at 5.6
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±0.1 when the pH0 of the solutions were 4.0, 5.6 and 8.5 probably due to the
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buffering ability of the catalyst.17 However, if the solution pH0 is very high or very
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low then the buffering ability of CuO-Fe3O4 becomes negligible. For example, in the
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case of 12.0 and 2.5, the solution pH did not change much during the reaction process.
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The loss of catalytic ability of CuO-Fe3O4 at pH0 2.5 is mainly due to the high
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leaching of the active component in the strong acid conditions. This can be seen from
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Figure 4(C) where values as high as 32.8 mg/L of copper ion and 0.6 mg/L of iron ion
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can be detected by ICP-OES after 120 min reaction. Another reason can be attributed
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to the adverse effect of H+ on PS activation.23, 25 In the pH0 range of 4.0 to 10.0, the
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leaching of copper ions gradually declined from 7.7 mg/L (pH0 4.0) to 1.3 mg/L (pH0
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10.0). By contrast, almost no leaching of iron ions can be detected at all these pH0
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values. The XRD characterization of the catalyst that was produced at different values
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of pH0 provides more powerful evidence, as shown in Figure S4. The typical peak of
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CuO at 2θ = 38.7° and 48.7° has totally disappeared at pH0 2.5 which indicates that
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this is responsible for the loss of catalytic ability of CuO-Fe3O4. This evidence also
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indicates that copper is the active component in the catalyst, whereas iron may not
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directly involve in the activation of PS.
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On the other hand, the rise of pH0 is also beneficial for the formation of surface
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hydroxyl groups on the catalyst, as illustrated by the increased intensity of the band at
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3454 cm-1 (Figure S5). The generated surface hydroxyl groups can behave as an
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active site for electron transfer and thus improve the performance of the catalyst.16, 26
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However, the enhanced removal of phenol in highly basic solution cannot be simply
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attributed to metal leaching and the formation of hydroxyl groups. The work
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investigated by Ahmad et al.27 could provide some insight into the good performance
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of CuO-Fe3O4/PS for the degradation of phenol at pH0 11.0 and 12.0. Based on their
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research, PS can be activated by the phenoxide ion which will be formed from phenol
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(pKa = 10.0) when pH0 is 11.0 and 12.0. Additionally, alkaline conditions will
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increase the reactively of PS and this could account for the increased removal of
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phenol at higher pH0.28 This is supported by control experiments, as shown in Figure
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S6, where 14% and 18% of phenol were removed at pH0 11.0 and 12.0 respectively, in
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the absence of any catalyst.
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Effect of NaHCO3. Bicarbonate is usually present in the aquatic environment at
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the concentration of 1.0 to 5.0 mM.29 Generally, bicarbonate was previously thought
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to have a negative effect on AOPs as it was a radical quencher.30 However, recently,
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some recent studies suggest that in both homogenous and heterogeneous AOPs,
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NaHCO3 can accelerate the degradation of pollutants.30, 31 As a preliminary step in
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investigating the potential application of CuO-Fe3O4/PS system in real water, the
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effect of bicarbonate on CuO-Fe3O4/PS system at its natural concentration was
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studied. As can be noticed from Figure 5(A), the removal of phenol increased with the
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addition of NaHCO3 and reached a plateau when 3.0 or 5.0 mM NaHCO3 were added.
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Interestingly, a further rise in NaHCO3 concentration to 10 mM slightly inhibited the
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degradation process but the amount removed (98%) is still about 18% higher than the
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control group (without NaHCO3). Analysis of the TOC results, as illustrated in Figure
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S1, showed a decrease in TOC of over 95% either at pH0 11.0 or in the presence of
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3.0 mM NaHCO3. This almost 100% removal of TOC means that the final
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degradation products of phenol were CO2 and H2O. This evidence once again
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suggests that the CuO-Fe3O4/PS system is highly effective for the treatment of organic
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wastewater, even in the presence of bicarbonate.
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To fully understand the role of NaHCO3, pH variation and the leached free metal
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ions were also monitored. In general, due to the buffering ability of NaHCO3, the pH
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stayed constant at 8.0±0.5 [Figure 5(B)] when 1.0 to 10 mM NaHCO3 was added
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although the more NaHCO3 added the higher the solution pH achieved.
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Correspondingly, the amount of Cu2+ leached dropped from 5.6 mg/L to less than 0.1
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mg/L when 1.0 mM NaHCO3 was added [Figure 5(C)]. Furthermore, when more than
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5.0 mM NaHCO3 was added, almost no copper ions could be determined in solution.
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Consequently, it is safe to say that the increased solution pH and the decreased
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leaching of copper ions are mainly responsible for the better performance of
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CuO-Fe3O4/PS/NaHCO3 system. Another role that NaHCO3 plays may be its
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complexing ability with Cu2+. Very recently, Chen et al.
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species under different bicarbonate concentration and found that Cu2+ (0.03 mM)
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exists mostly as CuCO3 when the concentration of bicarbonate was between 1.0 and
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10 mM. They also suggested that only CuCO3 was the main species responsible for
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the formation of the reactive species, Cu(III), which further led to the decolorization
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of acid orange 7 in their Cu2+-HCO3−-H2O2 system. The formation of high oxidation
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state metals in bicarbonate/H2O2 coupled process is not new. In fact, one intermediate
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in the mechanism that Lane et al.32 speculated for Mn(II)−catalyzed epoxidations of
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alkenes by bicarbonate/H2O2 is the formation of Mn(IV). As PS has a similar structure
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to H2O2, it was believed that PS could react with bicarbonate ions associated with the
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formation of peroxymonocarbonate via eq 3, which may further accelerate the
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conversion of Cu(II) to Cu(III)31 but probably does not rely on the direct reaction
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between Cu(II) and PS.
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HCO3− + S2O82− → HCO4− + SO4
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−
31
+ SO3−
calculated the Cu(II)
(3)
Mechanistic study. Two major alternative mechanisms have been proposed on
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copper activated persulfate: the free radical process15-17,
22
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process
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from copper to persulfate, as shown in equations 4-6. By contrast, the non-radical
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process assumed that the outer-sphere interaction between copper and PS makes PS
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more active to some organic pollutants.
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Cu2+ + PMS/H2O2 → Cu+ + SO5•−/O2•−
versus a non-radical
23
. In general, the free radical is formed as a result of the electron transfer
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Cu + + PMS/H 2 O 2 → Cu 2+ + SO 4 •− /HO •
(5)
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Cu2+ + PMS/PS/H2O2 → Cu3+ + SO4•−/HO•
(6)
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As stated above, it seems that alternative mechanisms are involved in different
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situations. Firstly, the mechanism based on equations (4) and (5) is ruled out in our
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system as PS has limited reduction ability. This can be further confirmed by our
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control studies as shown in Figure 3 where it can be seen that only when Cu2+/PS was
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kinetically activated by ascorbic acid, can phenol be degraded to any extent. This
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result indicates that Cu+ (generated from the reduction of Cu2+ by ascorbic acid) but
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not Cu2+ was effective for the activation of PS and this was in line with previous
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reports. For example, Liang et al.22 found that Cu2+ reacts much less with PS than
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with copper oxidate. Similarly, Liu et al.33 explored the destruction of propachlor with
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a Cu2+/PS system and although 70% of propachlor could be removed, the process took
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72 hours. The higher activity of ≡Cu(II) on the catalyst surface than soluble Cu2+ in
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aqueous solution is likely due to the higher E0Cu(III)/Cu(II) in solid phase than that in
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aqueous solution.16 It was reported that the reduction potential of Cu(III)/Cu(II) is to
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be 2.3 V in solid form and 1.57 V in the ionized phase.16, 34 Thus much published
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work as well as our own results indicate that the formation of Cu(I) and hence a Cu(I)
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based mechanism is not possible in our system.
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The second aspect that is worthy of discussion is whether radicals are produced or
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not. In order to check this, radical quenching experiments were first explored. It can
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be noticed from Figure 6(A) that the removal of phenol is only slightly (10%)
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decreased when 2.5 M (500 times the molar concentration of PS) methanol is added to
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the reaction. The first speculation is that a non-radical process may be involved and
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this is consistent with the findings of Zhang et al.23, who reported a CuO activated PS
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system is effective for the removal of 2,4-dichlorophenol and does not rely on the
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formation of sulfate radicals. To verify this speculation, the EPR experiments was
302
performed to detect any radical generated. The result clearly shows the exsitance of
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both SO4•− and HO• radicals, as noticed from Figure 6(B). Furthermore, when 2.5 M
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methanol was added to the solution, the intensity of DMPO adducts in CuO-Fe3O4/PS
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system decreased but the EPR signals were still observed. This indicates alcohols fail
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to capture surface-adsorbed radicals completely and consequently the addition of 2.5
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M mathanol only scavenged about 10% of phenol degradation.
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As identified above, surface radicals are probably responsible for the destruction of
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phenol. Figure S5 shows the ATR-FTIR spectral change of CuO-Fe3O4 at pH0 11.0
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and in prescence of 3.0 mM NaHCO3. As can be seen, there is a new small band
311
formed at 1125 cm-1 as well as the blue shift of the band at 1180 cm-1, indicating the
312
fomation of a complex at the surface of CuO-Fe3O4.16 It is thus inferred the
313
interaction between Cu(II) and PS leads to the formation of a weak bond at the
314
surface of the CuO-Fe3O4 (equations 7 and 8), accompanied by the generation of
315
Cu(III) and SO4•− radicals. The formed Cu(III) is unstable in the absence of strong
316
chealting anions33, 35 and could result in the formation of HO• via eq 9.36
317
≡Cu(II) + S2O82- → ≡Cu(II)--- O3SO2SO32-
(7)
318
≡Cu(II)--- O3SO2SO32- →≡Cu(III) + SO4•− + SO4−
(8)
319
≡Cu(III) + H2O → ≡Cu(II) + HO• + H+
(9)
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320
Meanwhile, as the fabricated catalyst is a homogeneous mixture of CuO and Fe3O4,
321
it is important to exam the role of each metal oxide. The control experiments using
322
CuO or Fe3O4 as catalyst were investigated. Figure S8 shows only 11% of phenol was
323
removed when PS was activated by 0.3 g/L Fe3O4, while the phenol removal was 18%
324
when Fe3O4 was replaced by 0.05 g/L CuO. It should be noted a 0.05 g/L dosage of
325
CuO was selected based on the equivalent metal content of 0.3 g/L CuO-Fe3O4. A
326
comparable removal of phenol could be achieved only when at a higher dosage of
327
CuO, i.e., 0.3 g/L. Furthermore, EPR studies clearly show the intensity of DMPO
328
adducts signals in CuO-Fe3O4/PS system was much stronger than the signals in CuO
329
or Fe3O4 activated PS system [see Figure 6(B)]. The synergistic effect of the
330
combined CuO-Fe3O4 catalyst may result from the reduction of Cu(III) by Fe(II) as
331
shown in eq 10, considering E0Cu(III)/Cu(II) = 2.3 V 16, 34 and E0Fe(III)/Fe(II) = 0.77 V 15.
332
≡Cu(III)
+
≡Fe(II) → ≡Fe(III) + ≡Cu(II)
(10)
333
To further verify this speculation, XPS spectra of fresh and used CuO-Fe3O4 were
334
recorded and the results were shown in Figure 7. Regarding the fresh catalyst, the
335
high-resolution spectra of the peaks at 710.6 ev and 713.0 ev are indicative of the
336
presence of Fe(II) and Fe(III)11. For the XPS spectra of Cu 2p3/2 region, the main peak
337
at binding energy of 933.6 ev is assigned to Cu(II)37 and the satelite peaks at 940.9 ev
338
and 943.2 ev are basically similar to that of dominated Cu(II) oxide species.38, 39 After
339
reaction, while copper species remains in Cu(II) status, the proportion of Fe(II)
340
species declined from 53.0% to 44.1%, indicating the oxidation of Fe(II) into Fe(III)
341
species. This is probabaly attributed to the reaction between Fe(II) and PS as well as
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the redox reaction between Cu(III) and Fe(II) (see eq 10).
343
On the basis of all the results obatained above, a possible mechanism for PS
344
activation by CuO-Fe3O4 is proposed. First of all, a Fenton-like reaction occured
345
between Cu(II) and S2O82− at the surface of CuO-Fe3O4 associated with the formation
346
of Cu(III) and SO4•−, which may lead to the formation of HO• radicals via equations 9
347
and 11, respectively. In particular, the regeneration of Cu(II) by Fe(II) favors the
348
continuous decomposition of PS as well as the production of radicals at the surface of
349
catalyst. In the meanwhile, the suface-adsorbed radicals may diffuse into the aqueous
350
solution. All these radicals as well as Cu(III) account for the destruction of phenol.
351
H2O + SO4•− → H+ + SO42− + HO•
(11)
352
Stability of CuO-Fe3O4. For the stability evaluation of catalyst, CuO-Fe3O4 was
353
magnetically separated and then washed 3 times with 1.5 L distilled water and dried at
354
105 °C. The stability of catalyst in two differnt reaction conditions (pH0 5.6 and pH0
355
11.0) was evaluated by their reuse performance. As can be seen from Figure 8 (A),
356
when the catalyst was reused at pH0 5.6, phenol removal decreased greatly from 80%
357
to less than 33% over 3 runs. This corresponds to the leaching of copper ions [Figure
358
2 and Figure 3(C)] as loss of Cu2+ will result in the decrease of active sites (Cu(II)) on
359
the catalyst surface. On the other hand, as phenol was not 100% mineralized at pH0
360
5.6 (Figure S1), any intermediate products may attach to catalyst surface so both the
361
leached Cu2+ and intermediates may result in the decrease in catalyst activity and thus
362
poor repeated catalyst performance. However, when the solution pH0 is 11.0, over
363
70% of phenol could still be removed in the 3rd run [ Figure 8(B)]. And it should also
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364
be noted that phenol could be further removed if the reaction time is extended beyond
365
120 min (data not shown). In order to identify the change in surface area, the BET
366
properties of the CuO-Fe3O4 after reaction were also measured (Table S1). After 3
367
runs, the BET surface area increased slightly to 31.63 and 34.02 m2/g at pH 5.6 and
368
pH 11.0, respectively and similar trend is observed for the pore size. The greater
369
surface area and pore size may due to the leaching of metals. However, the
370
CuO-Fe3O4 still remained in the spinel crystalline form in the cubic phase at pH 11.0
371
(Figure S4). All these information suggests that the CuO-Fe3O4 catalyst is relatively
372
stable. As very limited leached Cu2+ was detected by ICP-OES, the reason for the
373
gradually decline of the catalytic activity at pH 11.0 is not obvious and will be the
374
focus of future work.
375
Supporting Information Available
376
Supporting information associated with this article includes Table S1 and Figures
377
S1–S8. This material is available free of charge at http://pubs.acs.org.
378
Acknowledgement
379
This work is financially supported by the National Natural Science Foundation of
380
China (Grant No. 20977069). The generous help of Professor David H. Bremner in
381
polishing this manuscript is also greatly appreciated.
382
References
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on advanced oxidation processes for the removal of taste and odor compounds from aqueous media. Water Res. 2014, 53, 215-234. 4. Xiang, L.; Royer, S.; Zhang, H.; Tatibouet, J. M.; Barrault, J.; Valange, S., Properties of iron-based mesoporous silica for the CWPO of phenol: a comparison between impregnation and co-condensation routes. J. Hazard. Mater. 2009, 172, (2-3), 1175-11784. 5. Yao, Y.; Cai, Y.; Lu, F.; Qin, J.; Wei, F.; Xu, C.; Wang, S., Magnetic ZnFe2O4–C3N4 hybrid for photocatalytic degradation of aqueous organic pollutants by visible light. Ind. Eng. Chem. Res. 2014, 53, (44), 17294-17302. 6. Xu, L.; Wang, J., Fenton-like degradation of 2, 4-dichlorophenol using Fe3O4 magnetic nanoparticles. Appl. Catal. B: Environ. 2012, 123, 117-126. 7. Hou, L.; Zhang, Q.; Jérôme, F.; Duprez, D.; Zhang, H.; Royer, S., Shape-controlled nanostructured magnetite-type materials as highly efficient Fenton catalysts. Appl. Catal. B: Environ. 2014, 144, 739-749. 8. Pham, A. L.T.; Lee, C.; Doyle, F. M.; Sedlak, D. L., A silica-supported iron oxide catalyst capable of activating hydrogen peroxide at neutral pH values. Environ. Sci. Technol. 2009, 43, (23), 8930-8935. 9. Lei, Y.; Zhang, H.; Wang, J.; Ai, J., Rapid and continuous oxidation of organic contaminants with ascorbic acid and a modified ferric/persulfate system. Chem. Eng. J. 2015, 270, 73-79. 10. Luo, W.; Zhu, L.; Wang, N.; Tang, H.; Cao, M.; She, Y., Efficient removal of organic pollutants with magnetic nanoscaled BiFeO3 as a reusable heterogeneous Fenton-like catalyst. Environ. Sci. Technol. 2010, 44, (5), 1786-1791. 11. Xu, L.; Wang, J., Magnetic nanoscaled Fe3O4/CeO2 composite as an efficient Fenton-like heterogeneous catalyst for degradation of 4-chlorophenol. Environ. Sci. Technol. 2012, 46, (18), 10145-10153. 12. Yao, Y.; Yang, Z.; Zhang, D.; Peng, W.; Sun, H.; Wang, S., Magnetic CoFe2O4–graphene hybrids: facile synthesis, characterization, and catalytic properties. Ind. Eng. Chem. Res. 2012, 51, (17), 6044-6051. 13. Fu, Y.; Wang, X., Magnetically Separable ZnFe2O4–Graphene Catalyst and its High Photocatalytic Performance under Visible Light Irradiation. Ind. Eng. Chem. Res. 2011, 50, (12), 7210-7218. 14. Deng, J.; Shao, Y.; Gao, N.; Tan, C.; Zhou, S.; Hu, X., CoFe2O4 magnetic nanoparticles as a highly active heterogeneous catalyst of oxone for the degradation of diclofenac in water. . J. Hazard. Mater. 2013, 262, 836-844. 15. Ding, Y.; Zhu, L.; Wang, N.; Tang, H., Sulfate radicals induced degradation of tetrabromobisphenol A with nanoscaled magnetic CuFe2O4 as a heterogeneous catalyst of peroxymonosulfate. Appl. Catal. B: Environ. 2013, 129, 153-162. 16. Zhang, T.; Zhu, H.; Croue, J. P., Production of sulfate radical from peroxymonosulfate induced by a magnetically separable CuFe2O4 spinel in water: efficiency, stability, and mechanism. Environ. Sci. Technol. 2013, 47, (6), 2784-2791. 17. Guan, Y. H.; Ma, J.; Ren, Y. M.; Liu, Y. L.; Xiao, J. Y.; Lin, L. Q.; Zhang, C., Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals.
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Water Res. 2013, 47, (14), 5431-8. 18. Yao, Y.; Cai, Y.; Lu, F.; Wei, F.; Wang, X.; Wang, S., Magnetic recoverable MnFe2O4 and MnFe2O4-graphene hybrid as heterogeneous catalysts of peroxymonosulfate activation for efficient degradation of aqueous organic pollutants. J. Hazard. Mater. 2014, 270, 61-70. 19. Dorsey, A.; Ingerman, L.; Swarts, S., Toxicological profile for copper. Department of Health & Human Services, Public Health Service, Agency for Toxic Substances and Disease Registry: 2004. 20. Ahmed, M. M.; Chiron, S., Solar photo-Fenton like using persulphate for carbamazepine removal from domestic wastewater. Water Res. 2014, 48, 229-36. 21. Deng, Y.; Ezyske, C. M., Sulfate radical-advanced oxidation process (SR-AOP) for simultaneous removal of refractory organic contaminants and ammonia in landfill leachate. Water Res. 2011, 45, (18), 6189-94. 22. Liang, H.Y.; Zhang, Y.Q.; Huang, S.B.; Hussain, I., Oxidative degradation of p-chloroaniline by copper oxidate activated persulfate. Chem. Eng. J. 2013, 218, 384-391. 23. T. Zhang, Y. Chen, Y. Wang, J. Le Roux, Y. Yang, J.P. Croue, Efficient peroxydisulfate activation process not relying on sulfate radical generation for water pollutant degradation, Environ. Sci. Technol. 2014, 48, (10), 5868-5875. 24. Masomboon, N.; Ratanatamskul, C.; Lu, M.C., Chemical oxidation of 2, 6-dimethylaniline in the Fenton process. Environ. Sci. Technol. 2009, 43, (22), 8629-8634. 25. Guan, Y. H.; Ma, J.; Li, X. C.; Fang, J. Y.; Chen, L. W., Influence of pH on the formation of sulfate and hydroxyl radicals in the UV/peroxymonosulfate system. Environ. Sci. Technol. 2011, 45, (21), 9308-14. 26. Zhang, T.; Li, C.; Ma, J.; Tian, H.; Qiang, Z., Surface hydroxyl groups of synthetic α-FeOOH in promoting OH generation from aqueous ozone: property and activity relationship. Appl. Catal. B: Environ. 2008, 82, (1), 131-137. 27. Ahmad, M.; Teel, A. L.; Watts, R. J., Mechanism of persulfate activation by phenols. Environ. Sci. Technol. 2013, 47, (11), 5864-5871. 28. Furman, O. S.; Teel, A. L.; Watts, R. J., Mechanism of base activation of persulfate. Environ. Sci. Technol. 2010, 44, (16), 6423-6428. 29. Stiff, M., Copper/bicarbonate equilibria in solutions of bicarbonate ion at concentrations similar to those found in natural water. Water Res. 1971, 5, (5), 171-176. 30. Zhou, L.; Song, W.; Chen, Z.; Yin, G., Degradation of organic pollutants in wastewater by bicarbonate-activated hydrogen peroxide with a supported cobalt catalyst. Environ. Sci. Technol. 2013, 47, (8), 3833-3839. 31. Cheng, L.; Wei, M.; Huang, L.; Pan, F.; Xia, D.; Li, X.; Xu, A., Efficient H2O2 oxidation of organic dyes catalyzed by simple copper(II) ions in bicarbonate aqueous solution. Ind. Eng. Chem. Res. 2014, 53, (9), 3478-3485. 32. Lane, B. S.; Vogt, M.; DeRose, V. J.; Burgess, K., Manganese-catalyzed epoxidations of alkenes in bicarbonate solutions. J. Am. Chem. Soc. 2002, 124, (40), 11946-11954.
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477 478 479 480 481 482 483 484 485 486 487 488 489 490 491 492 493 494 495
33. Liu, C. S.; Shih, K.; Sun, C. X.; Wang, F., Oxidative degradation of propachlor by ferrous and copper ion activated persulfate. Sci. Total Environ. 2012, 416, 507-512. 34. Popova, T.; Aksenova, N., Complexes of copper in unstable oxidation states. Russian J. Coord. Chem. 2003, 29, (11), 743-765. 35. Chandra, S.; Yadava, K., Oxidation of some phenolic compounds with Cu (III). Microchem. J. 1970, 15, (1), 78-82. 36. Pham, A. N.; Xing, G.; Miller, C. J.; Waite, T. D., Fenton-like copper redox chemistry revisited: Hydrogen peroxide and superoxide mediation of copper-catalyzed oxidant production. J. Catal. 2013, 301, 54-64. 37. Espinos, J.; Morales, J.; Barranco, A.; Caballero, A.; Holgado, J.; González-Elipe, A., Interface effects for Cu, CuO, and Cu2O deposited on SiO2 and ZrO2. XPS determination of the valence state of copper in Cu/SiO2 and Cu/ZrO2 catalysts. J. Phys. Chem. B 2002, 106, (27), 6921-6929. 38. Moulder, J. F.; Chastain, J.; King, R. C., Handbook of X-ray photoelectron spectroscopy: a reference book of standard spectra for identification and interpretation of XPS data. Perkin-Elmer Eden Prairie, MN: 1992. 39. Li, T.T.; Cao, S.; Yang, C.; Chen, Y.; Lv, X.J.; Fu, W.F., Electrochemical water oxidation by in situ-generated copper oxide film from [Cu(TEOA)(H2O)2][SO4] Complex. Inorg. Chem. 2015, 54, (6), 3061-3067.
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(A)
Fabricated catalyst Fe3O4 (JCPDS: 65-3107)
Intensity /a.u.
CuO (JCPDS: 45-0937)
10
20
30
40
50
60
70
2-Theta (degree)
496
Figure 1. XRD pattern (A) and SEM-EDS data (B) of the fabricated catalyst.
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8
0.8
6
C/C0
0.6 4 0.4
CuO-Fe3O4 PS CuO-Fe3O4 + PS
0.2
2
Concentration (mg/L)
1.0
leached Fe leached Cu 0.0
0 0
20
40
60
80
100
120
Time (min) 497
Figure 2. Degradation efficiency of phenol and concentrations of leached metal ions.
498
Conditions: [PS] = 5.0 mM, [phenol] = 0.1 mM, [CuO-Fe3O4] = 0.3 g/L, pH0 5.6 ±
499
0.1.
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1.0
0.8
C/C0
0.6
0.4 5.6 mg/L Cu2+ 32 mg/L Cu2+ 64 mg/L Cu2+ 64 mg/L Cu2+ + 0.5 mM H2A
0.2
0.0 0
20
40
60
80
100
120
Time (min) 500
Figure 3. Degradation of phenol with Cu2+ /PS process. Conditions: [PS] = 5.0 mM,
501
[phenol] = 0.1 mM. [Cu2+] = 5.6 – 64 mg/L, pH0 5.6 ± 0.1, H2A means ascorbic acid.
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1.0 2.5 4.0 5.6 8.5 10.0 11.0 12.0
(A)
0.8
C/C0
0.6
0.4
0.2
0.0 0
20
40
60
80
100
120
Time (min)
12
(B)
pH value
10
8
2.5 5.6 10.0 12.0
4.0 8.5 11.0
60
80
6
4
2 0
20
40
100
Time (min)
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leaching concentration （ mg/L)
Page 27 of 32
(C)
Cu Fe
30 25 20 15 10 5 0
almost no metal leached
2.5
4.0
5.6
8.5
10.0
11.0
12.0
502
Figure 4 (A). Influence of pH0 on the degradation of phenol in CuO-Fe3O4/PS system.
503
(B) pH variation and (C) metal leaching during the reaction in the CuO-Fe3O4/PS
504
system. Conditions: [PS] = 5.0 mM, [phenol] = 0.1 mM, [CuO-Fe3O4] = 0.3 g/L,
505
Time = 120 min.
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1.0 0 0.1 1.0 3.0 5.0 10.0
0.8
C/C0
0.6
mM mM mM mM mM mM
(A)
0.4
0.2
0.0 0
20
40
60
80
100
120
Time (min)
9
pH value
8
(B) 7 0 mM 1.0 mM 5.0 mM
6
0.1 mM 3.0 mM 10.0 mM
5 0
20
40
60
80
100
Time (min)
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leaching concentration (mg/L)
7
(C)
Cu Fe
6 5 4 3 2
almost no metal leached
1 0
0 mM
0.1 mM
1.0 mM
3.0 mM
5.0 mM
10.0 mM
506
Figure 5 (A). Influence of NaHCO3 concentration on the degradation of phenol in the
507
CuO-Fe3O4/PS system. (B) pH variation and (C) metal leaching during the reaction in
508
the CuO-Fe3O4 /PS/NaHCO3 system. Conditions: [PS] = 5.0 mM, [phenol] = 0.1 mM,
509
[CuO-Fe3O4] = 0.3 g/L, Time = 120 min.
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1.0
(A)
C/C0 (phenol)
0.8
0.6
0.4
0.2
2.5 M methanol control
0.0 0
20
40
60
80
100
120
Time (min)
(B)
Intensity (a.u.)
Fe3O4/PS
♥ ♥
♣
♣
CuO-Fe3O4/PS
CuO/PS
DMPO-SO4
♣:
DMPO-OH
3200
3220
♣ ♥ ♥ ♥♣ ♣
♥
♥ ♥
♥
♥ ♥
♣
♥:
♣
♣♥
♥
♣
3240
♣ ♥ ♥
♣
♣ ♥
3260
♥
♥
3280
3300
3320
Magnetic field (G) 510
Figure 6. Effects of methanol on the degradation of phenol and in CuO-Fe3O4/PS
511
system (A) and EPR spectra in CuO-Fe3O4/PS, CuO/PS and Fe3O4 systems (B).
512
Conditions (A): [PS] = 5.0 mM, [phenol] = 0.1 mM, [CuO-Fe3O4] = 0.3 g/L,
513
Conditions (B): [PS] = 5.0 mM, [CuO] = 0.05 g/L, [Fe3O4] = 0.3 g/L, [DMPO] = 10
514
mM, pH0 5.6.
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515
(A)
Fe 2p3/2
Fe 2p1/2
fresh
Intensity/ a. u.
Fe(II) 53.0%
47.0% Fe(III)
used
Fe(II) 44.1%
55.9% Fe(III) 740
730
720
710
700
Binding energy (ev) 516 517
(B)
Cu 2p3/2
Intensity/ a. u.
fresh
satelite peaks
Cu(II) 100%
Cu(II) 100%
used
950
945
940
935
930
925
Binding energy (ev) 518
Figure 7．XPS spectra for Fe 2p regions (A) and Cu 2p regions (B) of fresh and used CuO-Fe3O4.
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1.0
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(A)
0.8
C/C0
0.6
0.4 1st run 2nd run 3rd run
0.2
0.0 0
60
120
180
240
300
360
Time (min)
1.0
(B)
0.8
C/C 0
0.6
0.4 1st run 2nd run 3rd run
0.2
0.0 0
60
120
180
240
300
360
Time (min) 519
Figure 8. Catalytic property of CuO-Fe3O4 for repeated use at pH0 5.6 (A) and pH0
520
11.0 (B). Conditions: [PS] = 5.0 mM, [phenol] = 0.1 mM, [CuO-Fe3O4] = 0.3 g/L.
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