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Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles
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Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles Huajuan Ling a , Kyungduk Kim a , Zongwen Liu a , Jeffrey Shi a , Xunjin Zhu b , Jun Huang a,∗ a b
Laboratory of Catalysis Engineering, School of Chemical and Biomolecular Engineering, The University of Sydney, Sydney, NSW 2006, Australia Department of Chemistry, Hong Kong Baptist University, Hong Kong
a r t i c l e
i n f o
Article history: Received 30 November 2014 Received in revised form 24 March 2015 Accepted 28 March 2015 Available online xxx Keywords: Photocatalysis Phenol degradation Titanium dioxide Hydrogen peroxide Hydrogenation Acid treatment
a b s t r a c t The TiO2 nanoparticles in the diameter of 10–23 nm prepared in this study offered high photocatalytic activity to continuously produce the higher concentration of hydroxyl radicals than that stoichiometric produced from the highly oxidizing agent H2 O2 in the water. The initiate phenol degradation rate on the TiO2 nanocatalyst was ca. 6 times higher than that in the phenol degradation only derived by H2 O2 . The addition of H2 O2 with TiO2 could enhance the initial concentration of hydroxyl radicals for the higher degradation rate. However, overloading H2 O2 with TiO2 could only slightly increase the degradation rate of phenol, and overloading TiO2 decreased the phenol degradation rate immediately. Further enhancement for the phenol degradation rate has been realized by surface modification of TiO2 via liquid acid treatment or hydrogenation. It did not change the bulk structure and the morphology/size of TiO2 , but strongly enhanced the photocatalytic performance for the phenol degradation. The formation of Lewis acid Ti3+ sites on the blue TiO2 surface via hydrogenation contributed the higher phenol degradation rate than Brønsted acid sites on acid-treated TiO2 . The preparation and regeneration of blue TiO2 avoids the utilization of corrosive liquid acids but offers higher photocatalytivity, which is promising for the water treatment. © 2015 Published by Elsevier B.V.
1. Introduction Phenol and phenolic derivatives are common contaminants in industrial wastewater that pose risks to the environment [1,2]. The conventional technologies for phenol removal to meet its safety discharge level which is in the range of 0.1–1.0 mg/L [3] include adsorption, coagulation, biological treatment [2–4], and enzyme oxidation [1,5]. However, these technologies do not actually destroy the organic pollutants and could generate secondary pollutants. In the last decades, heterogeneous photocatalysis employing transition metal semiconductors, in particular on titanium dioxide (TiO2 ) has been proposed as a low cost and low energy consumption for water treatment which ensures complete mineralization of organic contaminates without generating harmful by-products [6–8]. Generally, the photocatalysis is initiated when a photon with energy, h, equal to or greater than the band gap energy, Ebg , reaches the surface of the photocatalyst, thus resulting in the formation of electrons in the conduction band and positive
∗ Corresponding author. Tel.: +61 2 9351 748; fax: +61 2 9351 2854. E-mail address: [email protected] (J. Huang).
holes in the valence band [9]. The generated electrons and holes can recombine or they can transfer to different positions on the photocatalyst surface, then become trapped to form hydroxyl radicals (• OH) which are essential for highly efficient photocatalysis [10,11]. However, competition of the separation and recombination of photo-generated electrons and holes leads to process inefficiencies [4,10,12] and long exposure time for complete mineralization of phenol [6,13]. TiO2 photocatalysts are suitable for water treatment due to its high photochemical stability, nontoxicity and inexpensiveness [14,15]. For effective photocatalysts to overcome the above challenge, the structure and surface of TiO2 have been modified and tailored. Nahar et al. have prepared iron-doped TiO2 by using Degussa P25 TiO2 powder as the parent catalyst for degradation of phenol [16]. However, the leaching of metals from TiO2 surface would deactivate photocatalysts and make them hard to be regenerated. Carpio et al. have investigated the phenol degradation using TiO2 nanocrystals supported on activated carbon with 50% degradation in 4 h [17]. However, the significant amount of phenol absorbed on activated carbon took much longer time to reach target degradation level. Colon et al. [8] have investigated the photocatalytic oxidation of phenol over acidic pre-treated TiO2 prepared by
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Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048
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nitric, sulphuric and phosphoric acids [18]. However, waste acids from pre-treatment are harmful and corrosive pollutants for the environment. Recently, many researchers have reported that maximizing light harness by the modification of TiO2 and improving acidity of the semiconductor are feasible and cost-effective ways to improve the photocatalytic activity of TiO2 [18–22]. Recently, Mao et al. have produced high efficient disordered TiO2 nanocrystals for the photocatalytic splitting of water via the combination of adding impurities and hydrogenation treatment [19]. Wang et al. have reported that hydrogenated TiO2 nanoparticles introduce a surface disorder structure, extending the light absorption to the near infrared range (∼1200 nm) [22]. This modification process is clean, however, the production of black TiO2 requires the hydrogenation under the high pressure with the relatively high operation risk. In this work, effective TiO2 nanoparticles with a dominant anatase phase and a very small fraction of rutile phase has been prepared by a sol-gel method and tested for the photocatalytic degradation of phenol with the addition of hydrogen peroxide (H2 O2 ). The kinetics of heterogeneous photocatalytic processes was studied by investigating the effect of operating parameters. Moreover, the blue TiO2 nanoparticles have been developed under atmosphere pressure with improved charge separation and the photocatalytic activity. The efficiency of phenol degradation on this blue TiO2 was higher than that of acidic pre-treated TiO2 with sulphuric acid (H2 SO4 ) and hydrochloric acid (HCl). 2. Experimental 2.1. Catalyst preparation and characterization The TiO2 nanoparticles were prepared by a sol-gel method. Titanium (IV) isopropoxide (TTIP, >97%, Sigma–Aldrich) was firstly dissolved in ethanol solvent (> 99.5%, Sigma–Aldrich) and distilled water was added to the solution with a molar ratio of TTIP: ethanol: water = 1:10:2. Hydrochloric acid (12.5 mol% in water, molar ratio HCl:TTIP = 1:1) was added dropwise to adjust the pH under continuous stirring for the hydrolysis process of the mixture in 30 min to form sols. After aging for 24 h, the obtained gels were dried under 373 K overnight and then calcined at 773 K for 2 h to obtain TiO2 nanoparticles. The blue TiO2 nanoparticles were obtained after the hydrogenation of TiO2 in H2 flow (80 mL/min) at 673 K for 24 h. The acidic pre-treated TiO2 nanoparticles were produced by adding the parent TiO2 into an acid solution (0.05 g TiO2 per mL 0.5 M H2 SO4 or 1.0 M HCl) under sonication for 30 min and then dried overnight at 423 K for H2 SO4 and 323 K for HCl. The crystalline phases present in the TiO2 nanoparticles were analyzed by a Shimadzu X-Ray Diffractometer (XRD-6000) with CuK␣ radiation operated at 40 kV and 30 mA at scan range from 10 to 70 deg with continuous scanning mode at a rate of 2◦ min−1 . The surface area of the catalysts was measured by an Autosorb IQ-C system according to the N2 adsorption isotherm at 77 K. An amount of 50 mg of the samples was degassed at 423 K for 12 h under vacuum before the measurement. The morphology and size of solid catalysts of the samples were observed with scanning electron microscopy (SEM) which recorded on a FESEM, Zeiss Ultra+ and transmission electron microscopy (TEM) by a Philips CM120 BioFilter. The crystallographic structure of samples was imaged by a high resolution transmission electron microscopy (HRTEM) by JEOL 2200FS. 2.2. Photocatalytic activity test Photodegradation of phenol was carried out in a 1.0 L capacity cylindrical Pyrex-glass batch photoreactor. The proper amount of
Fig. 1. The XRD patterns of the as-prepared TiO2 and the further hydrogenated blue TiO2 samples.
the illuminated TiO2 nanoparticles was mixed with 150 ml of phenol solution with initial concentration of 55 ppm (0.58 mM). Prior to irradiation this solution was magnetically stirred in the dark for about 30 min to reach the adsorption equilibrium so that the loss of compound due to adsorption can be taken into account. A proper volume of H2 O2 was added to the solution at the initiation of irradiation and compressed air was purged into the solution to maintain aerobic condition. The light was provided by a 230-W high pressure mercury UV lamp which its strongest emission light is with wavelength of 254 nm. The distance from the UV lamp to the solution was 8 cm. All experiments were conducted at room temperature and about 3 ml of the aqueous solution was collected at regular intervals and analyzed by UV–vis spectroscopy at wavelength of 269.5 nm. The samples were filtered through a Millipore membrane filter with pore size of 0.1 m before the analysis. The Langmuir-Hinshelwood model is usually used to describe the kinetics of photocatalytic reactions of aquatic organics [7,23,24]: r=−
dC kr Kad C = 1 + Kad C dt
(1)
where kr is the intrinsic rate constant, Kad is the adsorption equilibrium constant and C is the concentration of aquatic organic. When the initial concentration of phenol is low, Eq. (1) can be simplified to the first-order kinetics with an apparent rate constant kapp [7]: ln
C C0
= −kr Kad t = −kapp t
(2)
The half-life time (degradation of phenol to its 50%) is calculated by the following equation: t1/2 =
ln 2 kapp
(3)
3. Results and discussion 3.1. TiO2 nanoparticles characterization The crystalline phases and crystallite size of both TiO2 and blue TiO2 sample were determined from the X-ray diffraction patterns (Fig. 1). The width of the peak obtained for both photocatalysts was narrow which implies that crystalline structure is presented
Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048
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Fig. 2. SEM image of (a) the as-prepared TiO2 , TEM images of (b) the as-prepared TiO2 , (c) S-TiO2 , (d) the hydrogenated blue TiO2 , HRTEM images of (e) the as-prepared TiO2 , (f) S-TiO2 , (g) the hydrogenated blue TiO2 .
Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048
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Fig. 3. Raman spectra of the as-prepared TiO2 , HCl-TiO2 , S-TiO2 and the hydrogenated blue TiO2 .
in both TiO2 samples. The crystalline phases present were mainly anatase (A) (1 0 1), (0 0 4), (1 0 3), (1 1 2), (2 0 0), (1 0 5), (2 1 1) and (2 0 4) spacing [25–29] and small portion of rutile (R) (1 1 0), (1 0 1) and (3 1 0) spacing [25–28,30,31], respectively. The broadening of the peak at the 25.023◦ was chosen to estimate the grain size of the sample which was approximately 21 nm by applying Scherrer’s formula [21]. No any change in XRD patterns were observed for TiO2 after acidic pre-treatment with H2 SO4 and HCl (not shown). After hydrogenation to transfer white TiO2 to blue one, however, the diffraction peak (1 0 1) of blue TiO2 (Fig. 1 right top) slightly shifted toward the higher angle, which indicates the reduction of the interplanar distance of anatase phase. This result clearly shows that the change of surface structure occurs during the hydrogenation process, but the bulk structure of particles kept the same. This can be ascribed to the escape of O atoms from the lattice on the surface, which often results in Ti3+ on surface. The morphology and particle size of all TiO2 catalysts were examined by SEM and TEM. The SEM image of TiO2 has demonstrated spherical shapes of the particles (Fig. 2(a)) and the TEM image shown in Fig. 2(b–d) has illuminated the circle shape crystalline structure of as-prepared, acid-pretreated, and blue TiO2 with a size of 10–23 nm that is generally in agreement with the size and main crystalline determined by XRD. Similarly, Raman spectra (Fig. 3) also showed the similar bulk structure of as-prepared, acid-pretreated, and blue TiO2 nanoparticles. There are five Ramanactive modes with frequencies at 143.7, 196.8, 395.5, 515.1 and 636.4 cm−1 presented Fig. 3. From the HRTEM images (Fig. 2(e–g)), the TiO2 nanoparticles are highly crystallized as seen from the well-resolved lattice, the acidic pre-treatment hardly changes the surface crystallinity, while the blue TiO2 nanoparticles had unique disordered surface structure after hydrogenation and the width of the disordered outer layer surrounding of the crystalline core was 1–2 nm. It was also confirmed by the slightly shift of the diffraction peak (1 0 1) during the formation of surface Ti3+ and the escape of O atoms from the lattice on the surface. 3.2. Optimize the reaction conditions on photodegradation of phenol 3.2.1. H2 O2 and catalyst addition for the generation of hydroxyl radicals In order to minimize the recombination of photo-generated electron-hole pairs and long exposure time for complete
Fig. 4. Phenol (55 pm, 0.58 mM) degradation rate as a function of time at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
mineralization of phenol, the integrated approach of TiO2 photocatalysis and oxidation process at the presence of H2 O2 has been attempted to improve the efficiency of photodegradation of phenol. As shown in Fig. 4, the direct photodegradation of phenol without TiO2 had extremely low phenol removal and only 3% phenol (Fig. 4a) was reduced after 24 h. It indicates UV itself is not able to degrade phenol in water. Adding H2 O2 for the advance UV promoted oxidation process, the degradation rate of phenol was obviously increased to 20% after 24 h (Fig. 4b). Phenol degradation by the advance oxidation process by H2 O2 involves the generation and subsequent reaction of hydroxyl radicals (• OH). The hydroxyl radical generated by hydrogen peroxide photolysis (Eq. (4)) oxidizes phenol from hydrogen abstraction to produce a phenoxyl radical (Eq. (5)) [3,32,33]. This organic radical quickly reacts with dissolved oxygen to yield peroxyl radicals, initiating thermal reactions of oxidative degradation, leading finally to carbon dioxide, water and inorganic salts (Eq. (6)) [3,32]. H2 O2 + h → 2• OH • OH + C H OH 6 5
C 6 H5
O•
→ C 6 H5
(4) O•
+ H2 O
+ O2 → Intermediate → CO2 + H2 O
(5) (6)
Based on the stoichiometry of the advance oxidation process by H2 O2 (Eq. (7)), 14 moles of H2 O2 are required for complete degradation of 1 mole phenol [34]. In this study, the initial phenol concentration was 0.58 mM and the concentration of H2 O2 was 4.82 mM, 60% of the stoichiometry amount was used and 20% degradation rate of phenol was obtained. Obviously, only adding the highly oxidizing agent could not complete the degradation of stable phenol in water. C6 H5 OH + 14H2 O2 → 6CO2 + 17H2 O
(7)
However, only introducing 0.5 g/L TiO2 photocatalyst strongly enhanced the reaction rate (without H2 O2 ), and achieved the total mineralization of phenol to CO2 and water in 23 h (Fig. 4c). It indicates that the TiO2 photocatalyst could continuously produce hydroxyl radicals for phenol removal. On TiO2 , oxygen as electron acceptor provides a natural sink for the photo-generated electrons. Thus, hydroxyl ions (OH− ) are traps for holes resulting in the formation of hydroxyl radicals (• OH) via Eq. (9) and (10) that are the major way of photo-generating • OH radicals in TiO2 photocatalysis process [35]. While the adsorbed oxygen species are traps for electrons, resulting in the formation of unstable superoxide species,
Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048
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Table 1 The kinetic parameters of different oxidation conditions on phenol degradation at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
Fig. 5. The degradation kinetic of phenol (55 pm, 0.58 mM) on different oxidation conditions at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
such as O2 •− and HO2 • (Eq. (11) and (12)) [10,36]. Hydrogen peroxide is produced via Eqs. (8)–(13) and further being reacted to form ·OH radicals (Eq. (4)) [35]. Phenol can be hydroxylated by the strongly oxidizing • OH radical that leads to successive oxidation and eventual ring opening [37]. The resulted intermediates will be further carboxylated to carbon dioxide and water (Eq. (5) and (6)) [37]. − + h+ TiO2 + h → eCB VB
(8)
h+ + OH− VB ads
(9)
→ • OHads
h+ + H2 Oads → • OHads + H+ VB •−
e− + O2,ads → O2 •−
•
+
O2 + H → HO2 •−
•
O2 + H+ + HO2 → H2 O2 + O2
(10) (11) (12) (13)
It was reported that the advance oxidation process by H2 O2 resulted in an initially fast photodegradation as high initial H2 O2 concentration, then followed by a slow degradation due to the loss in the effectiveness of hydroxyl radicals as a result of saturation [7]. Using TiO2 photocatalyst, hydroxyl radicals could be continuously produced and kept at a high level in the water, which contributed the strongly enhanced degradation rate of phenol. Further adding 4.82 mM of H2 O2 with 0.5 g/L TiO2 , the mineralization of phenol was completed in 15 h. Hence, the high initial hydroxyl radical concentration would promote the high degradation of phenol [39]. As the initial concentration of phenol is low, the reaction was proposed to obey the first-order kinetics with an apparent rate constant (kapp ). Plotting the negative of the natural logarithm of the fraction of unreacted phenol with respect to the irradiation time (for the oxidation conditions of TiO2 only and the addition of H2 O2 with TiO2 , the first 8 h reaction time was used) at different oxidation conditions for a first-order reaction gives a straight line with the slope of the line equals to kapp (Eq. (2), Fig. 5). The values of correlation factor R2 obtained from the linear regression are all above 0.93 which means these processes obey the first-order kinetics. The kap for phenol degradation under the four different oxidation conditions ranged from 0.0015 h−1 by the direct photodegradation of phenol without adding TiO2 and H2 O2 to 0.1169 h−1 using TiO2 and H2 O2 (Fig. 5), were used to estimate the initiate rate of phenol degradation (r0 ) and the half-life of these processes. The results are summarized in Table 1. The r0 of the direct photodegradation
Condition
kapp (h−1 )
R2
r0 (mM/h)
Half-life (h)
(a) Direct photolysis (b) UV/H2 O2 (c) UV/TiO2 (d) UV/TiO2 /H2 O2
0.0015 0.0101 0.0583 0.1169
0.9624 0.9358 0.9963 0.9965
0.0009 0.0059 0.0341 0.0679
462 69 12 6
of phenol without TiO2 , the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 and the addition of TiO2 with H2 O2 were 0.0009 mM/h, 0.0059 mM/h, 0.0341 mM/h and 0.0679 mM/h, respectively. The reaction rates of the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 only, and the photocatalysis on TiO2 with H2 O2 were increased by 7 folds, 38 folds and 76 folds, respectively, compared to the direct photodegradation of phenol by UV only. Moreover, the performance of the photocatalysis on TiO2 with H2 O2 was two times faster than the photocatalysis on TiO2 only. The half-life of the direct photodegradation of phenol by UV only, the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 and the photocatalysis on TiO2 with H2 O2 obtained from Eq. (3) was 462 h, 69 h, 12 h and 6 h, respectively. 3.2.2. Effect of TiO2 catalyst loading, initial phenol concentration, and hydrogen peroxide concentration Normally, high catalyst loading always promotes the reaction process. However, the optimal catalyst to phenol ratio is necessary to clarify due to the control of the operation costs. Fig. 6(a) clearly shows that increased in the amount of catalyst loading up to 1.0 g/L, the degradation rate of phenol increased, while further increasing in the amount of catalyst for the reaction, the degradation rate decreased. A number of studies have demonstrated that the rate of photodegradation for organic pollutants is strongly affected by the number of active sites and the photo-absorption ability of the catalyst used [7]. An increase in the amount of the catalyst increases the number of active sites for the generation of more radicals to high efficient of the process. However, excess amount of catalyst would result in negative effect as high suspension inhibits the penetration of the photon flux and the tendency toward agglomeration, resulting in a reduction in catalyst surface area available for light absorption and hence a drop in the photocatalytic degradation rate. The trade-off between these two opposing phenomena results in an optimum catalyst loading for the photocatalytic reaction which is 1.0 g/L. Previous studies have reported that the photocatalytic reaction rate of photocatalysis using TiO2 depends on the concentration of organic pollutants [12]. Fig. 6(b) shows that increased in the concentration of phenol, the degradation rate increased up to 0.36 mM and then slightly decreased with further increasing in phenol concentration. When the initial phenol concentration is extremely low, the number of surface activity sites of the fixed catalyst extremely sufficing for molecules of phenol, OH− , • OH and O2 •− adsorbed on the surface of the catalyst [3,12]. Increasing the initial concentration of phenol, the exposure time required for complete mineralization of phenol to CO2 and water increases. But the degradation rate increased up to a certain limited due to the total amount of phenol degraded actually increases. Further increasing the concentration of phenol, the degradation rate slightly decreased. This might be due to more and more molecules of phenol adsorbed on the surface of the catalyst, the competitive adsorption of OH− on the same site decreases and consequently the amount of • OH and O2 •− on the surface of catalyst decreases [3,12]. Furthermore, the formation of intermediates which are insoluble in water and
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Fig. 6. The optimization of operating parameters on phenol degradation. The degradation rate of catalyst loading (a), initial phenol concentrations (b) and hydrogen peroxide concentrations (c) on phenol degradation at room temperature.
probably stick on the surface of TiO2 lower the rate of phenol degradation at higher phenol concentration [37]. The slow reaction rates still remain a challenge in photodegradation of organic pollutants. As previously described, adding H2 O2 for the high initial concentration of hydroxyl radicals (• OH) could significantly improve the photocatalytic degradation rate of phenol in water. Fig. 6(c) shows that the degradation rate significantly accelerated when the concentration of H2 O2 was increased from 1.94 mM to 9.70 mM, and then slightly increased with further increasing in the concentration of H2 O2 . This is mainly due to the generation of more hydroxyl ions (OH− ) which are the likely traps for holes, leading to the formation of hydroxyl radicals (• OH). Further increasing the concentration of H2 O2 up to 48.3 mM and 77.6 mM, it will decrease the actual concentration of H2 O2 in the system as hydroxyl radical efficiently reacts with H2 O2 and produces HO2 • [40], also • OH radicals reacts with the generated HO2 • to produce water and oxygen [41]. Hence, there is a critical hydrogen peroxide concentration for high efficient of photocatalytic degradation of phenol. 3.3. The degradation performance of surface modified TiO2 After hydrogenation or acidic pre-treatment, the bulk crystal and morphological structure of TiO2 still maintained, evidenced by the XRD patterns, Raman spectra, and TEM images. However, the various functional active sites have been generated on the TiO2 surface. On blue TiO2 , the removal of O atoms during the hydrogenation would form the oxygen vacancies and active Ti3+ sites. The formation of Lewis acid Ti3+ sites on surface could narrow the band gap of the TiO2 nanoparticles and extend the light absorption
Fig. 7. Phenol (55 pm, 0.58 mM) degradation rate as a function of time at room temperature over the as-prepared TiO2 , the hydrogenated blue TiO2 , S-TiO2 and HCl-TiO2 .
toward the infrared range (∼1200 nm) [22]. For the acid-treated TiO2 , Brønsted acid sites were incorporated to the TiO2 surface to decomposition of the relatively stable side-on peroxide [42]. Both hydrogenation and acid-treatment of TiO2 surface could enhance the activity of catalysts in phenol degradation. As shown in Fig. 7,
Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048
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the degradation rate on the as-prepared TiO2 after 1 h reaction was 10.2%, while it was enhanced up to 24.6% on S-TiO2 , and 38.8% on HCl-TiO2 , and 44.2% on blue TiO2 , respectively. Among them, blue TiO2 performed the highest degradation of phenol. After reaction performed in 4 h, the degradation rates were continuously increasing on all samples such as 36.9% on the asprepared TiO2 , 44.8% on S-TiO2 , 49.9% on HCl-TiO2 , and 55.7% on the blue TiO2 , respectively. The photocatalytic activity of the S-TiO2 , HCl-TiO2 , and blue TiO2 nanoparticles were still higher than the as-prepared TiO2 , however, the difference between their catalytic performance became small. It may be caused by the leaching of Brønsted acid sites or the re-oxidation of Ti3+ sites. The enhancement of the photocatalytic activity of S-TiO2 and HCl-TiO2 nanoparticles causes by an excess of adsorbed protons (Brønsted acid sites) on the TiO2 surface after the acid treatment [18,42]. These protons do not dissolve in the organic solvent, but the leaching of protons from the surface was hard to avoid. Also, the oxygen vacancies on blue TiO2 surface would be slowly filled by O atoms from H2 O2 and Ti3+ sites could be transferred to Ti4+ sites accordingly. Compared with acid treatment of TiO2 , the preparation and regeneration of blue TiO2 avoids the utilization of corrosive liquid acids and the emission of waste water. The blue TiO2 provides a green way to enhance the photocatalytic activity of the photocatalysts in water treatment.
4. Conclusions One significant challenge for the degradation of organic compounds in water is to generate enough hydroxyl radicals to achieve the total mineralization of organic compounds to CO2 and water. The direct photodegradation of phenol under UV irradiation had extremely low phenol removal of 3% in 24 h (0.0009 mM/h). It indicates UV itself is not able to generate hydroxyl radicals to degrade phenol in water. Adding H2 O2 for the advance UV promoted oxidation process, the degradation rate of phenol was obviously increased to 20% after 24 h (0.0059 mM/h) due to the hydroxyl radical generated by hydrogen peroxide photolysis. However, 60% of the stoichiometry amount H2 O2 was used and only 20% degradation rate of phenol was obtained. Obviously, only adding the highly oxidizing agent could not complete the degradation of stable phenol in water. Using convenient sol-gel methods, this research prepared TiO2 fine nanoparticles in the diameter of 10–23 nm. Replacing H2 O2 by the TiO2 photocatalyst, the degradation rate of phenol was strongly enhanced and the total mineralization of phenol was achieved in 23 h (0.0341 mM/h). It indicates that the TiO2 photocatalyst could continuously produce high concentration of hydroxyl radicals for phenol removal. The addition of H2 O2 with TiO2 could enhance the initial concentration of hydroxyl radicals for the higher degradation rate (0.0679 mM/h). Overloading H2 O2 with TiO2 could only slightly increase the degradation rate of phenol, and overloading TiO2 decreased the phenol degradation rate immediately. Further enhancement for the phenol degradation rate can be realized by surface modification of TiO2 via liquid acid treatment or hydrogenation. The surface modification did not change the bulk structure and the morphology/size of TiO2 , but strongly enhanced the photocatalytic performance for the phenol degradation (24.6% on S-TiO2 , and 38.8% on HCl-TiO2 , and 44.2% on blue TiO2 , respectively, compared with 10.2% on as-prepared TiO2 ). The formation of Lewis acid Ti3+ sites on surface could narrow the band gap of the TiO2 nanoparticles, and Brønsted acid sites were incorporated to the TiO2 surface to decomposition of the relatively stable side-on peroxide. Among
7
them, blue TiO2 performed the highest degradation of phenol. After long-term running, the degradation rate on surface modified TiO2 increased slowly. It may be caused by the leaching of Brønsted acid sites or the re-oxidation of Ti3+ sites and could be regenerated. Compared with acid treatment of TiO2 , the preparation and regeneration of blue TiO2 avoiding the utilization of corrosive liquid acids but offering higher photocatalytivity is promising for the water treatment. Acknowledgment Financial support by The University of Sydney Early Career Research Scheme is gratefully acknowledged. References [1] A. Bódalo, J.L. Gómez, E. Gómez, J. Bastida, M.F. Máximo, Chemosphere 63 (2006) 626–632. [2] F. Shahrezaei, Y. Mansouri, A.A.L. Zinatizadeh, A. Akhbari, Powder Technol. 221 (2012) 203–212. [3] N. Kashif, F. Ouyang, J. Environ. Sci. (China) 21 (2009) 527–533. [4] S. Lathasree, A.N. Rao, B. SivaSankar, V. Sadasivam, K. Rengaraj, J. Mol. Catal. A: Chem. 223 (2004) 101–105. [5] J.L. Gómez, A. Bódalo, E. Gómez, A.M. Hidalgo, M. Gómez, M.D. Murcia, Chem. Eng. J. 145 (2008) 142–148. [6] S. Ahmed, M.G. Rasul, R. Brown, M.A. Hashib, J. Environ. Manage. 92 (2011) 311–330. [7] C.-H. Chiou, C.-Y. Wu, R.-S. Juang, Chem. Eng. J. 139 (2008) 322–329. [8] C.L. Wong, Y.N. Tan, A.R. Mohamed, J. Environ. Manage. 92 (1–2) (2011) 1669–1680, http://dx.doi.org/10.1016/j.jphotochem.2005.07.007 [9] S. Gupta, M. Tripathi, Chin. Sci. Bull. 56 (2011) 1639–1657. [10] H.d. Lasa, B. Serrano, M. Salaices, Photocatalytic Reaction Engineering, Springer, New York, 2005. [11] J. Liqiang, W. Dejun, W. Baiqi, L. Shudan, X. Baifu, F. Honggang, S. Jiazhong, J. Mol. Catal. A: Chem. 244 (2006) 193–200. [12] M.N. Chong, B. Jin, C.W.K. Chow, C. Saint, Water Res. 44 (2010) 2997–3027. [13] M.A. Gondal, M.N. Sayeed, Z. Seddigi, J. Hazard. Mater. 155 (2008) 83–89. [14] A. Mills, S. LeHunte, J. Photochem. Photobiol. A 108 (1997) 1–35. [15] G.H. Tian, H.G. Fu, L.Q. Jing, C.G. Tian, J. Hazard. Mater. 161 (2009) 1122–1130. [16] M.S. Nahar, K. Hasegawa, S. Kagaya, Chemosphere 65 (2006) 1976–1982. ˜ [17] E. Carpio, P. Zúniga, S. Ponce, J. Solis, J. Rodriguez, W. Estrada, J. Mol. Catal. A: Chem. 228 (2005) 293–298. ˜ M.C. Hidalgo, J.A. Navío, J. Photochem. Photo[18] G. Colón, J.M. Sánchez-Espana, biol. A: Chem. 179 (1–2) (2006) 20–27, http://dx.doi.org/10.1016/j.jphotochem. 2005.07.007 [19] X.B. Chen, L. Liu, P.Y. Yu, S.S. Mao, Science 331 (2011) 746–750. [20] Y.H. Hu, Angew. Chem. Int. Ed. 51 (2012) 12410–12412. [21] P. Saravanan, K. Pakshirajan, P. Saha, J. Hydro-environ. Res. 3 (2009) 45–50. [22] W. Wang, Y.R. Ni, C.H. Lu, Z.Z. Xu, RSC Adv. 2 (2012) 8286–8288. [23] R.W. Matthews, Water Res. 24 (1990) 653–660. [24] R.W. Matthews, J. Phys. Chem. 91 (1987) 3328–3333. [25] M. Asiltürk, F. Sayılkan, E. Arpac¸, J. Photochem. Photobiol. A 203 (2009) 64–71. [26] F. Li, L.-x. Guan, M.-l. Dai, J.-j. Feng, M.-m. Yao, Ceram. Int. 39 (2013) 7395–7400. [27] C. Wen, Y.-J. Zhu, T. Kanbara, H.-Z. Zhu, C.-F. Xiao, Desalination 249 (2009) 621–625. [28] Z. Wu, Z. Sheng, Y. Liu, H. Wang, N. Tang, J. Wang, J. Hazard. Mater. 164 (2009) 542–548. [29] C. Yu, Q. Fan, Y. Xie, J. Chen, Q. shu, J.C. Yu, J. Hazard. Mater. 237–238 (2012) 38–45. [30] A. Fujishima, X.T. Zhang, C.R. Chim. 9 (2006) 750–760. [31] J.-l. Li, X.-x. Ma, M.-r. Sun, X.-m. Li, Z.-l. Song, Trans. Nonferrous Met. Soc. 19 (Suppl. 3) (2009) s665–s668. [32] O. Legrini, E. Oliveros, A.M. Braun, Chem. Rev. 93 (1993) 671–698. [33] J. Bonin, I. Janik, D. Janik, D.M. Bartels, J. Phys. Chem. A 111 (2007) 1869–1878. [34] M.F. Kabir, E. Vaisman, C.H. Langford, A. Kantzas, Chem. Eng. J. 118 (2006) 207–212. [35] Y. Wang, C.-s. Hong, Water Res. 33 (1999) 2031–2036. [36] C.K. Grätzel, M. Jirousek, M. Grätzel, J. Mol. Catal. 60 (1990) 375–387. ´ ´ [37] A. Sobczynski, Ł. Duczmal, W. Zmudzinski, J. Mol. Catal. A: Chem. 213 (2004) 225–230. [39] R.M. Alberici, W.F. Jardim, Water Res. 28 (1994) 1845–1849. [40] N. Daneshvar, M.A. Behnajady, Y.Z. Asghar, J. Hazard. Mater. 139 (2007) 275–279. [41] S. Harimurti, A.U. Rahmah, A.A. Omar, T. Murugesan, J. Appl. Sci. 12 (2010) 1093–1099. [42] Q. Wang, M.A. Zhang, C.C. Chen, W.H. Ma, J.C. Zhao, Angew. Chem. Int. Ed. 49 (2010) 7976–7979.
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Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles Huajuan Ling a , Kyungduk Kim a , Zongwen Liu a , Jeffrey Shi a , Xunjin Zhu b , Jun Huang a,∗ a b
Laboratory of Catalysis Engineering, School of Chemical and Biomolecular Engineering, The University of Sydney, Sydney, NSW 2006, Australia Department of Chemistry, Hong Kong Baptist University, Hong Kong
a r t i c l e
i n f o
Article history: Received 30 November 2014 Received in revised form 24 March 2015 Accepted 28 March 2015 Available online xxx Keywords: Photocatalysis Phenol degradation Titanium dioxide Hydrogen peroxide Hydrogenation Acid treatment
a b s t r a c t The TiO2 nanoparticles in the diameter of 10–23 nm prepared in this study offered high photocatalytic activity to continuously produce the higher concentration of hydroxyl radicals than that stoichiometric produced from the highly oxidizing agent H2 O2 in the water. The initiate phenol degradation rate on the TiO2 nanocatalyst was ca. 6 times higher than that in the phenol degradation only derived by H2 O2 . The addition of H2 O2 with TiO2 could enhance the initial concentration of hydroxyl radicals for the higher degradation rate. However, overloading H2 O2 with TiO2 could only slightly increase the degradation rate of phenol, and overloading TiO2 decreased the phenol degradation rate immediately. Further enhancement for the phenol degradation rate has been realized by surface modification of TiO2 via liquid acid treatment or hydrogenation. It did not change the bulk structure and the morphology/size of TiO2 , but strongly enhanced the photocatalytic performance for the phenol degradation. The formation of Lewis acid Ti3+ sites on the blue TiO2 surface via hydrogenation contributed the higher phenol degradation rate than Brønsted acid sites on acid-treated TiO2 . The preparation and regeneration of blue TiO2 avoids the utilization of corrosive liquid acids but offers higher photocatalytivity, which is promising for the water treatment. © 2015 Published by Elsevier B.V.
1. Introduction Phenol and phenolic derivatives are common contaminants in industrial wastewater that pose risks to the environment [1,2]. The conventional technologies for phenol removal to meet its safety discharge level which is in the range of 0.1–1.0 mg/L [3] include adsorption, coagulation, biological treatment [2–4], and enzyme oxidation [1,5]. However, these technologies do not actually destroy the organic pollutants and could generate secondary pollutants. In the last decades, heterogeneous photocatalysis employing transition metal semiconductors, in particular on titanium dioxide (TiO2 ) has been proposed as a low cost and low energy consumption for water treatment which ensures complete mineralization of organic contaminates without generating harmful by-products [6–8]. Generally, the photocatalysis is initiated when a photon with energy, h, equal to or greater than the band gap energy, Ebg , reaches the surface of the photocatalyst, thus resulting in the formation of electrons in the conduction band and positive
∗ Corresponding author. Tel.: +61 2 9351 748; fax: +61 2 9351 2854. E-mail address: [email protected] (J. Huang).
holes in the valence band [9]. The generated electrons and holes can recombine or they can transfer to different positions on the photocatalyst surface, then become trapped to form hydroxyl radicals (• OH) which are essential for highly efficient photocatalysis [10,11]. However, competition of the separation and recombination of photo-generated electrons and holes leads to process inefficiencies [4,10,12] and long exposure time for complete mineralization of phenol [6,13]. TiO2 photocatalysts are suitable for water treatment due to its high photochemical stability, nontoxicity and inexpensiveness [14,15]. For effective photocatalysts to overcome the above challenge, the structure and surface of TiO2 have been modified and tailored. Nahar et al. have prepared iron-doped TiO2 by using Degussa P25 TiO2 powder as the parent catalyst for degradation of phenol [16]. However, the leaching of metals from TiO2 surface would deactivate photocatalysts and make them hard to be regenerated. Carpio et al. have investigated the phenol degradation using TiO2 nanocrystals supported on activated carbon with 50% degradation in 4 h [17]. However, the significant amount of phenol absorbed on activated carbon took much longer time to reach target degradation level. Colon et al. [8] have investigated the photocatalytic oxidation of phenol over acidic pre-treated TiO2 prepared by
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nitric, sulphuric and phosphoric acids [18]. However, waste acids from pre-treatment are harmful and corrosive pollutants for the environment. Recently, many researchers have reported that maximizing light harness by the modification of TiO2 and improving acidity of the semiconductor are feasible and cost-effective ways to improve the photocatalytic activity of TiO2 [18–22]. Recently, Mao et al. have produced high efficient disordered TiO2 nanocrystals for the photocatalytic splitting of water via the combination of adding impurities and hydrogenation treatment [19]. Wang et al. have reported that hydrogenated TiO2 nanoparticles introduce a surface disorder structure, extending the light absorption to the near infrared range (∼1200 nm) [22]. This modification process is clean, however, the production of black TiO2 requires the hydrogenation under the high pressure with the relatively high operation risk. In this work, effective TiO2 nanoparticles with a dominant anatase phase and a very small fraction of rutile phase has been prepared by a sol-gel method and tested for the photocatalytic degradation of phenol with the addition of hydrogen peroxide (H2 O2 ). The kinetics of heterogeneous photocatalytic processes was studied by investigating the effect of operating parameters. Moreover, the blue TiO2 nanoparticles have been developed under atmosphere pressure with improved charge separation and the photocatalytic activity. The efficiency of phenol degradation on this blue TiO2 was higher than that of acidic pre-treated TiO2 with sulphuric acid (H2 SO4 ) and hydrochloric acid (HCl). 2. Experimental 2.1. Catalyst preparation and characterization The TiO2 nanoparticles were prepared by a sol-gel method. Titanium (IV) isopropoxide (TTIP, >97%, Sigma–Aldrich) was firstly dissolved in ethanol solvent (> 99.5%, Sigma–Aldrich) and distilled water was added to the solution with a molar ratio of TTIP: ethanol: water = 1:10:2. Hydrochloric acid (12.5 mol% in water, molar ratio HCl:TTIP = 1:1) was added dropwise to adjust the pH under continuous stirring for the hydrolysis process of the mixture in 30 min to form sols. After aging for 24 h, the obtained gels were dried under 373 K overnight and then calcined at 773 K for 2 h to obtain TiO2 nanoparticles. The blue TiO2 nanoparticles were obtained after the hydrogenation of TiO2 in H2 flow (80 mL/min) at 673 K for 24 h. The acidic pre-treated TiO2 nanoparticles were produced by adding the parent TiO2 into an acid solution (0.05 g TiO2 per mL 0.5 M H2 SO4 or 1.0 M HCl) under sonication for 30 min and then dried overnight at 423 K for H2 SO4 and 323 K for HCl. The crystalline phases present in the TiO2 nanoparticles were analyzed by a Shimadzu X-Ray Diffractometer (XRD-6000) with CuK␣ radiation operated at 40 kV and 30 mA at scan range from 10 to 70 deg with continuous scanning mode at a rate of 2◦ min−1 . The surface area of the catalysts was measured by an Autosorb IQ-C system according to the N2 adsorption isotherm at 77 K. An amount of 50 mg of the samples was degassed at 423 K for 12 h under vacuum before the measurement. The morphology and size of solid catalysts of the samples were observed with scanning electron microscopy (SEM) which recorded on a FESEM, Zeiss Ultra+ and transmission electron microscopy (TEM) by a Philips CM120 BioFilter. The crystallographic structure of samples was imaged by a high resolution transmission electron microscopy (HRTEM) by JEOL 2200FS. 2.2. Photocatalytic activity test Photodegradation of phenol was carried out in a 1.0 L capacity cylindrical Pyrex-glass batch photoreactor. The proper amount of
Fig. 1. The XRD patterns of the as-prepared TiO2 and the further hydrogenated blue TiO2 samples.
the illuminated TiO2 nanoparticles was mixed with 150 ml of phenol solution with initial concentration of 55 ppm (0.58 mM). Prior to irradiation this solution was magnetically stirred in the dark for about 30 min to reach the adsorption equilibrium so that the loss of compound due to adsorption can be taken into account. A proper volume of H2 O2 was added to the solution at the initiation of irradiation and compressed air was purged into the solution to maintain aerobic condition. The light was provided by a 230-W high pressure mercury UV lamp which its strongest emission light is with wavelength of 254 nm. The distance from the UV lamp to the solution was 8 cm. All experiments were conducted at room temperature and about 3 ml of the aqueous solution was collected at regular intervals and analyzed by UV–vis spectroscopy at wavelength of 269.5 nm. The samples were filtered through a Millipore membrane filter with pore size of 0.1 m before the analysis. The Langmuir-Hinshelwood model is usually used to describe the kinetics of photocatalytic reactions of aquatic organics [7,23,24]: r=−
dC kr Kad C = 1 + Kad C dt
(1)
where kr is the intrinsic rate constant, Kad is the adsorption equilibrium constant and C is the concentration of aquatic organic. When the initial concentration of phenol is low, Eq. (1) can be simplified to the first-order kinetics with an apparent rate constant kapp [7]: ln
C C0
= −kr Kad t = −kapp t
(2)
The half-life time (degradation of phenol to its 50%) is calculated by the following equation: t1/2 =
ln 2 kapp
(3)
3. Results and discussion 3.1. TiO2 nanoparticles characterization The crystalline phases and crystallite size of both TiO2 and blue TiO2 sample were determined from the X-ray diffraction patterns (Fig. 1). The width of the peak obtained for both photocatalysts was narrow which implies that crystalline structure is presented
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Fig. 2. SEM image of (a) the as-prepared TiO2 , TEM images of (b) the as-prepared TiO2 , (c) S-TiO2 , (d) the hydrogenated blue TiO2 , HRTEM images of (e) the as-prepared TiO2 , (f) S-TiO2 , (g) the hydrogenated blue TiO2 .
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Fig. 3. Raman spectra of the as-prepared TiO2 , HCl-TiO2 , S-TiO2 and the hydrogenated blue TiO2 .
in both TiO2 samples. The crystalline phases present were mainly anatase (A) (1 0 1), (0 0 4), (1 0 3), (1 1 2), (2 0 0), (1 0 5), (2 1 1) and (2 0 4) spacing [25–29] and small portion of rutile (R) (1 1 0), (1 0 1) and (3 1 0) spacing [25–28,30,31], respectively. The broadening of the peak at the 25.023◦ was chosen to estimate the grain size of the sample which was approximately 21 nm by applying Scherrer’s formula [21]. No any change in XRD patterns were observed for TiO2 after acidic pre-treatment with H2 SO4 and HCl (not shown). After hydrogenation to transfer white TiO2 to blue one, however, the diffraction peak (1 0 1) of blue TiO2 (Fig. 1 right top) slightly shifted toward the higher angle, which indicates the reduction of the interplanar distance of anatase phase. This result clearly shows that the change of surface structure occurs during the hydrogenation process, but the bulk structure of particles kept the same. This can be ascribed to the escape of O atoms from the lattice on the surface, which often results in Ti3+ on surface. The morphology and particle size of all TiO2 catalysts were examined by SEM and TEM. The SEM image of TiO2 has demonstrated spherical shapes of the particles (Fig. 2(a)) and the TEM image shown in Fig. 2(b–d) has illuminated the circle shape crystalline structure of as-prepared, acid-pretreated, and blue TiO2 with a size of 10–23 nm that is generally in agreement with the size and main crystalline determined by XRD. Similarly, Raman spectra (Fig. 3) also showed the similar bulk structure of as-prepared, acid-pretreated, and blue TiO2 nanoparticles. There are five Ramanactive modes with frequencies at 143.7, 196.8, 395.5, 515.1 and 636.4 cm−1 presented Fig. 3. From the HRTEM images (Fig. 2(e–g)), the TiO2 nanoparticles are highly crystallized as seen from the well-resolved lattice, the acidic pre-treatment hardly changes the surface crystallinity, while the blue TiO2 nanoparticles had unique disordered surface structure after hydrogenation and the width of the disordered outer layer surrounding of the crystalline core was 1–2 nm. It was also confirmed by the slightly shift of the diffraction peak (1 0 1) during the formation of surface Ti3+ and the escape of O atoms from the lattice on the surface. 3.2. Optimize the reaction conditions on photodegradation of phenol 3.2.1. H2 O2 and catalyst addition for the generation of hydroxyl radicals In order to minimize the recombination of photo-generated electron-hole pairs and long exposure time for complete
Fig. 4. Phenol (55 pm, 0.58 mM) degradation rate as a function of time at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
mineralization of phenol, the integrated approach of TiO2 photocatalysis and oxidation process at the presence of H2 O2 has been attempted to improve the efficiency of photodegradation of phenol. As shown in Fig. 4, the direct photodegradation of phenol without TiO2 had extremely low phenol removal and only 3% phenol (Fig. 4a) was reduced after 24 h. It indicates UV itself is not able to degrade phenol in water. Adding H2 O2 for the advance UV promoted oxidation process, the degradation rate of phenol was obviously increased to 20% after 24 h (Fig. 4b). Phenol degradation by the advance oxidation process by H2 O2 involves the generation and subsequent reaction of hydroxyl radicals (• OH). The hydroxyl radical generated by hydrogen peroxide photolysis (Eq. (4)) oxidizes phenol from hydrogen abstraction to produce a phenoxyl radical (Eq. (5)) [3,32,33]. This organic radical quickly reacts with dissolved oxygen to yield peroxyl radicals, initiating thermal reactions of oxidative degradation, leading finally to carbon dioxide, water and inorganic salts (Eq. (6)) [3,32]. H2 O2 + h → 2• OH • OH + C H OH 6 5
C 6 H5
O•
→ C 6 H5
(4) O•
+ H2 O
+ O2 → Intermediate → CO2 + H2 O
(5) (6)
Based on the stoichiometry of the advance oxidation process by H2 O2 (Eq. (7)), 14 moles of H2 O2 are required for complete degradation of 1 mole phenol [34]. In this study, the initial phenol concentration was 0.58 mM and the concentration of H2 O2 was 4.82 mM, 60% of the stoichiometry amount was used and 20% degradation rate of phenol was obtained. Obviously, only adding the highly oxidizing agent could not complete the degradation of stable phenol in water. C6 H5 OH + 14H2 O2 → 6CO2 + 17H2 O
(7)
However, only introducing 0.5 g/L TiO2 photocatalyst strongly enhanced the reaction rate (without H2 O2 ), and achieved the total mineralization of phenol to CO2 and water in 23 h (Fig. 4c). It indicates that the TiO2 photocatalyst could continuously produce hydroxyl radicals for phenol removal. On TiO2 , oxygen as electron acceptor provides a natural sink for the photo-generated electrons. Thus, hydroxyl ions (OH− ) are traps for holes resulting in the formation of hydroxyl radicals (• OH) via Eq. (9) and (10) that are the major way of photo-generating • OH radicals in TiO2 photocatalysis process [35]. While the adsorbed oxygen species are traps for electrons, resulting in the formation of unstable superoxide species,
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Table 1 The kinetic parameters of different oxidation conditions on phenol degradation at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
Fig. 5. The degradation kinetic of phenol (55 pm, 0.58 mM) on different oxidation conditions at room temperature. (a) UV irradiation only, (b) UV irradiation with 4.82 mM H2 O2 , (c) UV irradiation with 0.5 g/L TiO2 , and (d) UV irradiation with 0.5 g/L TiO2 and 4.82 mM H2 O2.
such as O2 •− and HO2 • (Eq. (11) and (12)) [10,36]. Hydrogen peroxide is produced via Eqs. (8)–(13) and further being reacted to form ·OH radicals (Eq. (4)) [35]. Phenol can be hydroxylated by the strongly oxidizing • OH radical that leads to successive oxidation and eventual ring opening [37]. The resulted intermediates will be further carboxylated to carbon dioxide and water (Eq. (5) and (6)) [37]. − + h+ TiO2 + h → eCB VB
(8)
h+ + OH− VB ads
(9)
→ • OHads
h+ + H2 Oads → • OHads + H+ VB •−
e− + O2,ads → O2 •−
•
+
O2 + H → HO2 •−
•
O2 + H+ + HO2 → H2 O2 + O2
(10) (11) (12) (13)
It was reported that the advance oxidation process by H2 O2 resulted in an initially fast photodegradation as high initial H2 O2 concentration, then followed by a slow degradation due to the loss in the effectiveness of hydroxyl radicals as a result of saturation [7]. Using TiO2 photocatalyst, hydroxyl radicals could be continuously produced and kept at a high level in the water, which contributed the strongly enhanced degradation rate of phenol. Further adding 4.82 mM of H2 O2 with 0.5 g/L TiO2 , the mineralization of phenol was completed in 15 h. Hence, the high initial hydroxyl radical concentration would promote the high degradation of phenol [39]. As the initial concentration of phenol is low, the reaction was proposed to obey the first-order kinetics with an apparent rate constant (kapp ). Plotting the negative of the natural logarithm of the fraction of unreacted phenol with respect to the irradiation time (for the oxidation conditions of TiO2 only and the addition of H2 O2 with TiO2 , the first 8 h reaction time was used) at different oxidation conditions for a first-order reaction gives a straight line with the slope of the line equals to kapp (Eq. (2), Fig. 5). The values of correlation factor R2 obtained from the linear regression are all above 0.93 which means these processes obey the first-order kinetics. The kap for phenol degradation under the four different oxidation conditions ranged from 0.0015 h−1 by the direct photodegradation of phenol without adding TiO2 and H2 O2 to 0.1169 h−1 using TiO2 and H2 O2 (Fig. 5), were used to estimate the initiate rate of phenol degradation (r0 ) and the half-life of these processes. The results are summarized in Table 1. The r0 of the direct photodegradation
Condition
kapp (h−1 )
R2
r0 (mM/h)
Half-life (h)
(a) Direct photolysis (b) UV/H2 O2 (c) UV/TiO2 (d) UV/TiO2 /H2 O2
0.0015 0.0101 0.0583 0.1169
0.9624 0.9358 0.9963 0.9965
0.0009 0.0059 0.0341 0.0679
462 69 12 6
of phenol without TiO2 , the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 and the addition of TiO2 with H2 O2 were 0.0009 mM/h, 0.0059 mM/h, 0.0341 mM/h and 0.0679 mM/h, respectively. The reaction rates of the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 only, and the photocatalysis on TiO2 with H2 O2 were increased by 7 folds, 38 folds and 76 folds, respectively, compared to the direct photodegradation of phenol by UV only. Moreover, the performance of the photocatalysis on TiO2 with H2 O2 was two times faster than the photocatalysis on TiO2 only. The half-life of the direct photodegradation of phenol by UV only, the advanced oxidation process by H2 O2 , the photocatalysis on TiO2 and the photocatalysis on TiO2 with H2 O2 obtained from Eq. (3) was 462 h, 69 h, 12 h and 6 h, respectively. 3.2.2. Effect of TiO2 catalyst loading, initial phenol concentration, and hydrogen peroxide concentration Normally, high catalyst loading always promotes the reaction process. However, the optimal catalyst to phenol ratio is necessary to clarify due to the control of the operation costs. Fig. 6(a) clearly shows that increased in the amount of catalyst loading up to 1.0 g/L, the degradation rate of phenol increased, while further increasing in the amount of catalyst for the reaction, the degradation rate decreased. A number of studies have demonstrated that the rate of photodegradation for organic pollutants is strongly affected by the number of active sites and the photo-absorption ability of the catalyst used [7]. An increase in the amount of the catalyst increases the number of active sites for the generation of more radicals to high efficient of the process. However, excess amount of catalyst would result in negative effect as high suspension inhibits the penetration of the photon flux and the tendency toward agglomeration, resulting in a reduction in catalyst surface area available for light absorption and hence a drop in the photocatalytic degradation rate. The trade-off between these two opposing phenomena results in an optimum catalyst loading for the photocatalytic reaction which is 1.0 g/L. Previous studies have reported that the photocatalytic reaction rate of photocatalysis using TiO2 depends on the concentration of organic pollutants [12]. Fig. 6(b) shows that increased in the concentration of phenol, the degradation rate increased up to 0.36 mM and then slightly decreased with further increasing in phenol concentration. When the initial phenol concentration is extremely low, the number of surface activity sites of the fixed catalyst extremely sufficing for molecules of phenol, OH− , • OH and O2 •− adsorbed on the surface of the catalyst [3,12]. Increasing the initial concentration of phenol, the exposure time required for complete mineralization of phenol to CO2 and water increases. But the degradation rate increased up to a certain limited due to the total amount of phenol degraded actually increases. Further increasing the concentration of phenol, the degradation rate slightly decreased. This might be due to more and more molecules of phenol adsorbed on the surface of the catalyst, the competitive adsorption of OH− on the same site decreases and consequently the amount of • OH and O2 •− on the surface of catalyst decreases [3,12]. Furthermore, the formation of intermediates which are insoluble in water and
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Fig. 6. The optimization of operating parameters on phenol degradation. The degradation rate of catalyst loading (a), initial phenol concentrations (b) and hydrogen peroxide concentrations (c) on phenol degradation at room temperature.
probably stick on the surface of TiO2 lower the rate of phenol degradation at higher phenol concentration [37]. The slow reaction rates still remain a challenge in photodegradation of organic pollutants. As previously described, adding H2 O2 for the high initial concentration of hydroxyl radicals (• OH) could significantly improve the photocatalytic degradation rate of phenol in water. Fig. 6(c) shows that the degradation rate significantly accelerated when the concentration of H2 O2 was increased from 1.94 mM to 9.70 mM, and then slightly increased with further increasing in the concentration of H2 O2 . This is mainly due to the generation of more hydroxyl ions (OH− ) which are the likely traps for holes, leading to the formation of hydroxyl radicals (• OH). Further increasing the concentration of H2 O2 up to 48.3 mM and 77.6 mM, it will decrease the actual concentration of H2 O2 in the system as hydroxyl radical efficiently reacts with H2 O2 and produces HO2 • [40], also • OH radicals reacts with the generated HO2 • to produce water and oxygen [41]. Hence, there is a critical hydrogen peroxide concentration for high efficient of photocatalytic degradation of phenol. 3.3. The degradation performance of surface modified TiO2 After hydrogenation or acidic pre-treatment, the bulk crystal and morphological structure of TiO2 still maintained, evidenced by the XRD patterns, Raman spectra, and TEM images. However, the various functional active sites have been generated on the TiO2 surface. On blue TiO2 , the removal of O atoms during the hydrogenation would form the oxygen vacancies and active Ti3+ sites. The formation of Lewis acid Ti3+ sites on surface could narrow the band gap of the TiO2 nanoparticles and extend the light absorption
Fig. 7. Phenol (55 pm, 0.58 mM) degradation rate as a function of time at room temperature over the as-prepared TiO2 , the hydrogenated blue TiO2 , S-TiO2 and HCl-TiO2 .
toward the infrared range (∼1200 nm) [22]. For the acid-treated TiO2 , Brønsted acid sites were incorporated to the TiO2 surface to decomposition of the relatively stable side-on peroxide [42]. Both hydrogenation and acid-treatment of TiO2 surface could enhance the activity of catalysts in phenol degradation. As shown in Fig. 7,
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the degradation rate on the as-prepared TiO2 after 1 h reaction was 10.2%, while it was enhanced up to 24.6% on S-TiO2 , and 38.8% on HCl-TiO2 , and 44.2% on blue TiO2 , respectively. Among them, blue TiO2 performed the highest degradation of phenol. After reaction performed in 4 h, the degradation rates were continuously increasing on all samples such as 36.9% on the asprepared TiO2 , 44.8% on S-TiO2 , 49.9% on HCl-TiO2 , and 55.7% on the blue TiO2 , respectively. The photocatalytic activity of the S-TiO2 , HCl-TiO2 , and blue TiO2 nanoparticles were still higher than the as-prepared TiO2 , however, the difference between their catalytic performance became small. It may be caused by the leaching of Brønsted acid sites or the re-oxidation of Ti3+ sites. The enhancement of the photocatalytic activity of S-TiO2 and HCl-TiO2 nanoparticles causes by an excess of adsorbed protons (Brønsted acid sites) on the TiO2 surface after the acid treatment [18,42]. These protons do not dissolve in the organic solvent, but the leaching of protons from the surface was hard to avoid. Also, the oxygen vacancies on blue TiO2 surface would be slowly filled by O atoms from H2 O2 and Ti3+ sites could be transferred to Ti4+ sites accordingly. Compared with acid treatment of TiO2 , the preparation and regeneration of blue TiO2 avoids the utilization of corrosive liquid acids and the emission of waste water. The blue TiO2 provides a green way to enhance the photocatalytic activity of the photocatalysts in water treatment.
4. Conclusions One significant challenge for the degradation of organic compounds in water is to generate enough hydroxyl radicals to achieve the total mineralization of organic compounds to CO2 and water. The direct photodegradation of phenol under UV irradiation had extremely low phenol removal of 3% in 24 h (0.0009 mM/h). It indicates UV itself is not able to generate hydroxyl radicals to degrade phenol in water. Adding H2 O2 for the advance UV promoted oxidation process, the degradation rate of phenol was obviously increased to 20% after 24 h (0.0059 mM/h) due to the hydroxyl radical generated by hydrogen peroxide photolysis. However, 60% of the stoichiometry amount H2 O2 was used and only 20% degradation rate of phenol was obtained. Obviously, only adding the highly oxidizing agent could not complete the degradation of stable phenol in water. Using convenient sol-gel methods, this research prepared TiO2 fine nanoparticles in the diameter of 10–23 nm. Replacing H2 O2 by the TiO2 photocatalyst, the degradation rate of phenol was strongly enhanced and the total mineralization of phenol was achieved in 23 h (0.0341 mM/h). It indicates that the TiO2 photocatalyst could continuously produce high concentration of hydroxyl radicals for phenol removal. The addition of H2 O2 with TiO2 could enhance the initial concentration of hydroxyl radicals for the higher degradation rate (0.0679 mM/h). Overloading H2 O2 with TiO2 could only slightly increase the degradation rate of phenol, and overloading TiO2 decreased the phenol degradation rate immediately. Further enhancement for the phenol degradation rate can be realized by surface modification of TiO2 via liquid acid treatment or hydrogenation. The surface modification did not change the bulk structure and the morphology/size of TiO2 , but strongly enhanced the photocatalytic performance for the phenol degradation (24.6% on S-TiO2 , and 38.8% on HCl-TiO2 , and 44.2% on blue TiO2 , respectively, compared with 10.2% on as-prepared TiO2 ). The formation of Lewis acid Ti3+ sites on surface could narrow the band gap of the TiO2 nanoparticles, and Brønsted acid sites were incorporated to the TiO2 surface to decomposition of the relatively stable side-on peroxide. Among
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Please cite this article in press as: H. Ling, et al., Photocatalytic degradation of phenol in water on as-prepared and surface modified TiO2 nanoparticles, Catal. Today (2015), http://dx.doi.org/10.1016/j.cattod.2015.03.048