ZrO2 nanoparticles for dimethyl sulfide oxidation

ZrO2 nanoparticles for dimethyl sulfide oxidation

Journal of Colloid and Interface Science 446 (2015) 226–236 Contents lists available at ScienceDirect Journal of Colloid and Interface Science www.e...

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Journal of Colloid and Interface Science 446 (2015) 226–236

Contents lists available at ScienceDirect

Journal of Colloid and Interface Science www.elsevier.com/locate/jcis

Catalytic activity of Fe/ZrO2 nanoparticles for dimethyl sulfide oxidation Keshav Chand Soni ⇑, S. Chandra Shekar, Beer Singh, T. Gopi PD Division, Defence Research and Development Establishment, Gwalior 474002, India

g r a p h i c a l a b s t r a c t

a r t i c l e

i n f o

Article history: Received 18 November 2014 Accepted 14 January 2015 Available online 24 January 2015 Keywords: Ozonation Catalysis Oxidation Fe Nano TPR TPD

a b s t r a c t A low-temperature vapor phase catalytic oxidation of dimethyl sulfide (DMS) with ozone over nano-sized Fe2O3–ZrO2 catalyst is carried out at temperatures of 50–200 °C. Nanostructured Fe2O3–ZrO2 catalyst (FZN) is prepared by modified sol–gel method using citric acid as a chelating agent and conventional FZ catalyst is prepared with co-precipitation method. The catalysts are characterized using N2-BET surface area and pore size distributions, X-ray diffraction, TPR, TPD of DMS and NH3, SEM and TEM. The effects of operating temperature, ozone/DMS concentration and gas hourly space velocity (GHSV) on DMS removal efficiencies via catalytic ozonation are investigated. Relatively higher amount of ozone decomposition is observed on nanocatalyst compared to the co-precipitate catalyst from 50 °C to 150 °C. In contrast, at 200 °C irrespective of the particle size, both catalysts performed similar activity. It clearly demonstrates that under ozone assisted catalytic oxidation over nanocatalyst offers the 100% of DMS conversion at lower temperature. The synthesized nanocatalyst and ozone are observed highly efficient for low temperature catalytic oxidation of DMS. The stability test shows that the nanocatalyst have relatively high activity and stability under the reaction conditions. A plausible reaction mechanism has been proposed for the oxidation of DMS based on the possible reaction products. Ó 2015 Elsevier Inc. All rights reserved.

1. Introduction Volatile organic compounds (VOC), whether directly or indirectly, are known to be major contributors to air pollution. Volatile organic sulfur compounds (VOSC) are odorous, which are characterized by their high toxicity, potential corrosive effect, and very low odor threshold values (OTV) [1]. Dimethyl sulfide (DMS) is a typical gaseous odor pollutant owing to its offensive smell with a very low OTV of 0.6–40 ppb, and extreme negative hedonic characteristics [2]. It is generated in the off-gas from paper

⇑ Corresponding author. E-mail address: [email protected] (K.C. Soni). http://dx.doi.org/10.1016/j.jcis.2015.01.031 0021-9797/Ó 2015 Elsevier Inc. All rights reserved.

mills, wastewater treatment, kraft process, pharmaceuticals, insecticides and fungicides. Due to the potential odor pollutant, the treatment of DMS has attracted more and more attention. There have been numerous reported methods for the removal of VOC, including catalytic oxidation, thermal oxidation, and adsorption processes. Among these methods, catalytic oxidation is a promising effluent treatment method to control the emission of VOSC. The catalytic oxidation of reduced sulfur compounds with oxygen over metal oxide catalysts has shown high activity over a temperature range of 200–400 °C [3–5]. Due to the poisoning effect of sulfur on the catalysts, especially at lower temperatures (<200 °C), cost effective treatment technologies are needed to minimize the VOSC emission [5,6]. The use of ozone as an alternative oxidant for the catalytic oxidation process is receiving considerable

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attention not only from further reduction of operating temperature but also from the decrease of the activation energy compared to oxygen. [7–9]. A previous study has pointed out that ozone decomposition over a metal oxide catalyst plays an important role in oxidation process and manganese oxide has the highest activity for ozone decomposition among the p-type transition metal oxides [9]. It has been reported that complete methanol oxidation is achieved at 150 °C over V2O5/TiO2 catalyst with an ozone/methanol ratio of 1.2 and space velocity of 60,000 h 1 [10]. The activation energy for catalytic oxidation of chlorobenzene with ozone over iron oxide nanocatalyst is reported to be 20 kJ mol 1 (60–210 °C) while using oxygen is 43 kJ mol 1 (>200 °C) [8]. Xi et al. [11] reported that the addition of ozone significantly reduced the oxidative temperature of acetone on MnOx/Al2O3. The partial/total oxidation of methane to methanol and formaldehyde over Li/MgO catalysts, methanol to methylformate on supported vanadia catalysts is described in literature using ozone as oxidant [12,13]. Most of the low temperature gas-phase catalytic ozonation is focused on hydrocarbons, methanol, ethanol, acetone, chlorobenzene, and benzene [8,10–13]. A very little literature has been reported about the gas-phase catalytic ozonation for complete mineralization of reduced VOSC using transition metal oxide catalysts [14,15]. Recently, supported metal oxides nanoparticles have gained considerable attention due to their novel properties compared to their bulk materials [16,17]. Small particle size, high surface area, and more importantly, densely populated surface coordination unsaturated sites, could potentially provide improved catalytic performance of nanocrystalline transition metal oxides over respective bulk materials. The study of Iron oxides nanoparticles has attracted intensive attention over the past decades due to the potential applications in catalysis [18,19]. Recent studies showed nanosized iron particles are widely appearing to be economically feasible on an industrial scale for Fischer–Tropsch process [20–22]. In other report immobilized Iron (III)–substituted polyoxotungstates is found effective catalysts for solvent-free aerobic oxidation of n-hexadecane [23]. It is also reported that Fe nanocatalyst showed higher ethylbenzene, cyclohexene, and benzyl alcohol conversion (>84%) in the temperatures ranges of 50–120 °C and reaction time of 24 h [24]. Many industrially important reactions such as hydroprocessing, oxidation of alcohols and synthesis of methanol and higher alcohols, are reported to show high activity and selectivity using zirconia instead of other conventional supports such as alumina, silica, and titania [25–27]. Zirconia is a suitable support for iron oxide which develops unusually strongly acidic properties, very interesting surface characteristics and in many cases better selectivity and activity than the corresponding metal oxides supported on alumina and silica [25–28]. Okamoto [29] reported that Fe/ZrO2 catalysts exhibit high catalytic activity in the NO–CO reaction at low Fe content (<5 wt.%). The advantages of using ZrO2 as a catalytic support is its special chemical properties such as: (i) interacting strongly with the active phase; (ii) possessing a high thermal stability and being more chemically inert than classical oxides; (iii) being the only metal oxide which may possess all four chemical properties: namely acidity, basicity, reducing ability and oxidizing ability [30]. Several methods have been developed to prepare Iron oxide nanoparticles, such as co-precipitation [7], microemulsion [31], surfactant mediated or template synthesis [32], sol–gel processes [33], and so on. Sol–gel procedure is a promising method which provides excellent control of physical, chemical and microstructural properties at the molecular level [21,22,34]. Navio et al. [35] prepared Fe/ZrO2 samples by a sol–gel technique and observed that the introduction of a small amount of iron in ZrO2 produces a phase transition from monoclinic to tetragonal. They found that iron stabilizes cubic ZrO2 when introduced as Fe2O3.


The configuration involves a very short Fe–Zr distance that suggests the presence of a direct Fe–Zr metal to metal bond. The structure, morphology and size of ZrO2 supported nanoparticles are under the influence of various parameters such as pH, temperature, reaction time, reagent concentrations, catalyst nature, concentration, H2O/precursor molar ratio, chelating agent, aging temperature, time and drying. In this study, a modified sol–gel method is used to prepare Iron oxides nanoparticles, which avoids use of expensive alkoxides and involve mixing of metal precursor solution with an organic polyfunctional acid. In the multidentate chelated metal ions solution, ethylene glycol is added, and upon heating, polyesterification occurs. This leads to a homogeneous and viscous gel in which the metal ions are uniformly distributed. It is well-known that the citric acid is a good chelating agent and a suitable precursor that ensures high specific surface area and nanometric particle size [36,37]. The purpose of this study is to compare the vapor phase catalytic activity of Fe2O3–ZrO2 nanocatalyst with co-precipitate catalyst for the oxidation of DMS in the presence of ozone. This study highlights the effects of preparation method as well as process parameters governing the conversion and product selectivity of DMS for the long-term effectiveness of catalytic oxidation by ozone. 2. Materials and methods 2.1. Preparation of catalyst Fe/ZrO2 nanocatalyst was synthesized by a sol–gel method using citric acid as chelating agent. In a typical synthesis, 4.7 g of zirconium isopropoxide (Aldrich) was dissolved in 40 ml of isopropanol (Aldrich). A specific amount of ferric nitrate (Merck) was dissolved in 10 ml of isopropanol and added to the above solution. This solution was added under constant stirring at 70 °C to obtain a clear and homogeneous solution. Monohydrate citric acid (Aldrich) as a chelating agent was used in this solution and to regulate the pH, nitric acid was employed as a catalytic agent. Depending on the pH, nitric acid was mixed with ethylene glycol in accordance with the weight of 10% citric acid added glycol and introduced to the above solution to hydrolyze the zirconium isopropoxide. In this study, the molar ratios Fe/Zr were maintained at 12.5 and citric acid/Fe at 1.9. The solution pH was observed 1.8 in hydrolysis and condensation step. The obtained mixtures was further stirred for 30 min and then aged at 70 °C for 4 h to dry until a thick transparent gel being formed. The resulting gel was dried at 110 °C in oven for 12 h. The dried solid gel was ground to a fine powder, subsequently subjected to decomposition in N2 at 300 °C for 1 h and calcined in air for 3 h at 350 °C. The nanocatalyst was designated as FZN, containing 8 wt.% of Fe content. Another catalyst prepared by co-precipitation method. Requisite amount of zirconium oxychloride (Aldrich) and Feric chloride (Aldrich) were dissolved in 50 ml of deionised water and precipitated the respective hydroxides with an aqueous solution of 6% NH3 at pH 9.9. The precipitate was filtered and washed to make it free from chloride ion. The washed product was dried at 120 °C in air for 12 h followed by calcination in air for 3 h at 350 °C. The convectional catalyst containing 8 wt.% of Fe was designated as FZ catalyst. 2.2. Physicochemical characterization of catalysts Brunauer–Emmett–Teller (BET) and Barrett–Joyner–Halenda (BJH) methods were used to estimate the specific surface area and pore size distributions, respectively. Nitrogen adsorption– desorption isotherms were obtained using a surface area analyzer (Micromeritics ASAP-2010) at liquid N2 temperature ( 196 °C).


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Prior to analyses, 0.2 gm of the catalyst samples were loaded into a quartz reactor and degassed at 120 °C with a helium purge for 60 min. A five-point nitrogen adsorption isotherm was used to determine the BET surface area of the samples. X-ray diffraction (XRD) patterns of calcined and used forms of the catalysts are recorded on a Rigaku Miniflex (M/s. Rigaku Corporation, Japan) X-ray diffractometer using Ni filtered Cu Ka radiation (k = 1.5406 Å) with a scan speed of 2° min 1 and a scan range of 2–80° at 30 kV and 50 mA. Temperature programmed reduction (TPR) and Temperature programmed desorption (TPD) experiments were conducted on TPR/TPD Nuchrom unit equipped with a thermal conductivity detector. The TPR studies of the catalyst samples (0.2 g) were carried out using a 5.6% H2–Ar mixture (50 cc/min) and in the temperature range of 50–700 °C at a heating rate of 10 °C/min. In the NH3-TPD experiments, the catalysts were pre-treated in helium for 1 h and then NH3 pulses were dosed at 100 °C in order to saturate catalyst surface. The samples were then allowed to cool at room temperature (30 °C) and once the base line restored, the temperature was increased from 100 °C to 500 °C with a heating rate of 10 °C min 1 in flowing helium (50 cc/min). For thermal desorption experiments, DMS was allowed to adsorb on (0.2 g) of catalyst at 30 °C under comparable reaction conditions without using any oxidizer for 1 h and charged into the reactor for thermal desorption studies. TPD was performed by heating the sample from 40 °C to 500 °C at a rate of 10 °C min 1 using helium as carrier. The surface morphology of Fe supported catalysts was carried out on a scanning electron microscope (SEM, JEOL JAX840) operating at 20 kV. Micrographs were taken after coating by gold sputtering. Samples for transmission emission spectroscopy (TEM, JEM-200FX electron microscope) were prepared by keeping a drop of the colloidal solution prepared in acetone as solvent on a copper grid, coated with a thin carbon film. Samples were dried and kept in vacuum oven in a desiccator. TEM was operated between 160 kV and 180 kV. Ferret diameter was used in the analysis of particle size and its distribution. 2.3. Reaction studies for the catalytic oxidation Catalytic oxidation of DMS was carried out in a continuous flow, fixed-bed reactor (5 mm ID  150 mm length) at atmospheric pressure. The reactor was mounted vertically in an electrically heated tubular furnace to measure the temperature of the heating zone of the furnace. Another thermocouple was placed inside the reactor in contact with the catalyst bed to measure the temperature of catalyst bed. Prior to the experiments, the catalyst was calcined at 350 °C in air (100 ml/min) for an hour. Nitrogen saturated DMS (80 ppm DMS in N2; Aldrich AR 99%) gas stream at 25 °C was used as DMS feed. The DMS saturated vapors were passed through the reactor. The ozone stream was directly diffused onto the catalyst bed to minimize the gas phase reaction, which was kept at a pre-set temperature (50–200 °C). A temperature rise was observed in the catalyst bed as soon as the gas mixture was passed through the bed due to the exothermic nature of the oxidation reaction. This was controlled by an external cooling of the reactor bed by a strong flow of air over the reactor surface which maintained the constant reaction temperature. The reaction was studied in the temperature range of 50 °C to120 °C, with gas hourly space velocity from 12,000 h 1 to 21,000 h 1 and the ozone to DMS mole ratio were varied between 0.2 and 1.4. In the experiment the oxygen was used through ozonator and the experimental results were obtained by switch off the ozonator discharge power supply. Compressed O2 cylinder (99.7% dried over silica gel) was used for the catalytic reactions as well as for the ozone generator. The ozone generator was standardized by the iodometric method [38]. The flow rates of oxygen through the ozone generator, nitrogen, and DMS in N2 were adjusted using mass flow controllers (Sierra

Instruments Inc., flow accuracy ±1%). Analysis of the gaseous mixture feed and that exiting from the catalyst bed were carried out online using calibrated CO and CO2 analyzers (Technovation Instruments CO-MP94 and CO2-P89), which were capable of measuring concentrations of CO and CO2, respectively. The product mixture of condensable phase was trapped in acetonitrile solvent at 5 °C. The condensable phase was analyzed by GC equipped with PFPD/FID detectors. Qualitative analysis was done by GC– MS equipped with mass-selective detector. The GC–MS analyses reveal that SO2 and carbon oxide are the major products and small amount of partial oxidation products, such as and dimethyl sulfoxide (DMSO) and dimethyl sulfone (DMSO2) were formed. Trace amount of other product were also observed (<1% by area), which are not identified/quantified in present study. The conversion is defined as the number of moles of product formed per mole of the reactant per gram of the catalyst. The conversion of DMS was calculated from the amount of DMS in the inlet and outlet of the reactor, i. e., {([DMS]in [DMS]out)/[DMS]in}. After each test, the regeneration of catalysts was carried out in flowing ozone at 120 °C for 1 h. 3. Result and discussions 3.1. BET-surface area and pore size distribution The surface areas of the FZ and FZN calcined catalysts are 212 m2 g 1 and 247 m2 g 1, respectively. The BET specific surface area results for the modified sol–gel catalyst are higher than those observed for co-precipitated catalysts. When the metal component is loaded to the support, some micropores are blocked, causing the decreased specific surface area. When the catalyst is prepared via citric acid modification, the decrease in the surface area is not as much as in traditionally prepared catalyst. This discrepancy is attributed to the smaller size of the metal component particles on the ZrO2 surface, and to the fewer micropores blocked by the metal component. The pore size distribution obtained from N2 BET surface area results are displayed in Fig. 1. The catalyst pore size distribution reveals that the distribution in narrow range and most of the area is in the mesoporous range which is ascribed to the ZrO2 support. In FZN catalyst, two types pore size distributions observed in the range of 20–80 Å and besides small amount of micropore. The narrow pore size distribution is also observed in the range of 15–45 Å for FZ catalyst. It is observed that there is only one peak centered at 3.8 nm for FZ catalyst suggesting the presence of a uniform pore size. Furthermore, the peak shape is very narrow, indicating that the mesoporous ZrO2 have a sharply defined pore size distribution. It clearly indicates that the catalyst

Fig. 1. Pore size distribution of FZ and FZN catalyst.

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prepared by modified sol–gel method have mesopore as well as micropore with almost uniform pore size. It is also reported that higher surface areas and larger pore diameters (between 3 nm and 10 nm) is particularly useful for overcoming mass transfer limitations of reactant molecules and facilitate the reaction by amplifying the reaction surface [39–41]. 3.2. X-ray diffraction X-ray diffraction patterns of Fe/ZrO2 calcined catalysts are displayed in Fig. 2a. The FZ catalyst clearly indicates the presence of tetragonal ZrO2 phase with d values: 2.95, 2.54, 1.79 and 1.53. The diffraction signals at 2h = 30.2°, 50.7° and 60.2° are indexed as the ZrO2 reflection planes of (1 0 1), (2 0 0) and (2 1 1), respectively [CAS No. 79-1769]. The XRD pattern of FZ catalyst indicates the characteristic diffraction peaks at 2h values of 24.2°, 33.1°, 35.5° and 49.6° which correspond to a-Fe2O3 phase (JCPDS No. 33-0664). In the XRD pattern of FZN nanocatalyst, no separate FexOy as well as ZrO2 phase are observed which indicate the formation of Fe2O3–ZrO2 solid solution or very high dispersion of Fe on ZrO2 support. However, the possibility cannot be ruled out for the presence of Fe crystallites having size less than 5 nm, which is beyond the detection capacity of the powder X-ray diffraction technique. The XRD results of nanocatalyst represent the amorphous nature of Iron oxide and ZrO2 which retained under synthetic conditions whereas crystalline nature in co-precipitate catalyst [42]. Santos et al. [43] is reported that the nanostructured ZrO2 prepared by modified sol–gel method to be X–ray amorphous when calcined below 400 °C. 3.3. Temperature programmed reduction The reduction behavior of the FZ and FZN catalysts are studied by H2–TPR and the results are illustrated in Fig. 2b. It can be observed that the reduction process in H2 occurs in two distinct stages in the temperature range of 300–750 °C, as commonly observed in most of the iron-based catalysts. The first stage is ascribed to the transformations of Fe2O3 to Fe3O4, whereas the second stage represents the transformation of Fe3O4 to Fe. The first stage can be further divided into two peaks: the first peak corresponds to the reduction of Fe2O3 to Fe3O4, and the second

Fig. 2. (a) X-ray diffraction patterns of FZ and FZN catalysts, (b) temperature programmed reduction profiles of FZ and FZN catalysts.


peak corresponds to subsequent reductions of Fe3O4 to FeO. The reduction temperature for both catalysts is in line with the observations of those reported in literature [44–46]. On the other hand the reduction maxima is shifted to low temperature in both the reduction peaks, and the significant shift in reduction temperature is observed for Fe3O4 to Fe; approximately 100 °C compared to the corresponding co-precipitate catalyst. It is reported that enhanced interaction between iron and copper oxides nanocatalyst shifts the Fe2O3 to Fe3O4 peak to lower temperature in H2-TPR in comparison with other catalyst [45]. However, smaller crystal size in nanocatalyst has weak influence on the reduction of Fe3O4 to Fe, since thermodynamics and the nucleation of new crystal structures control the reduction rates at higher temperatures instead of H2 dissociation steps reported for Zn, Cu, and K promoted Fe based catalysts [47]. On the other hand It is also reported that the second peak can be shifted to lower temperatures for Fe2O3/SiO2; approximately 80 °C which indicated that the metal support interactions also influences the reduction maxima for Fe catalysts. It is concluded that the amount of catalyst reduction is increased by with decreasing the particle size from co-precipitate catalyst to nano size [21,48]. 3.4. Temperature programmed desorption of DMS In order to ascertain the desorption behavior of DMS on particle size effect, the thermal desorption of DMS is performed for the FZ and FZN catalysts (Fig. 3a) and the maximum desorption temperature and semi quantitative results are summarized in Table 1. The small peak centered at 160 °C is ascribed to desorption of weakly adsorbed DMS in both catalysts. The peak centered at 342 °C with onset temperature of 250 °C in case of FZN nanocatalyst is ascribed to the chemisorbed DMS. Similarly the peak centered at 396 °C with lower intensity ascribed to the chemisorbed DMS in case of FZ catalyst. The desorption peak at low temperature corresponds to the catalyst activity that indicates the maximum activity in ozone at 150 °C. However, the high temperature desorption peak corresponds to the thermal decomposition of DMS on the Fe sites as well as on ZrO2 acid sites. The relative desorption intensity of the high temperature peak for FZN nanocatalyst is higher than that of FZ catalyst. It appears that the populations of Fe nanoparticles as well as the ZrO2 acid sites are much higher in FZN nanocatalyst

Fig. 3. (a) Temperature programmed desorption profile of DMS over FZ and FZN catalysts, (b) temperature programmed desorption profile of ammonia over FZ and FZN catalysts.


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Table 1 Catalytic properties of Fe catalysts. Catalyst code

Fe (wt%)

Fe (wt% EDAX)

BET-SA (m2 g


8 8

6.74 6.52

212 247



than that of corresponding co-precipitated catalyst, which resulted in more amount of adsorption of DMS on FZN nanocatalyst. The XRD data obtained for the FZN catalyst also confirms the amorphous nature of ZrO2 [43]. These results also confirm the observations for adsorption of DMS on various Fe2O3 active sites as well as ZrO2 acid sites [7]. The TPD profile of DMS confirms the significant variation in the semi quantitative analysis of the active sites and acid sites observed in this study. The strong adsorption of DMSO on acid sites might be responsible for extended oxidations which lead to formation of COx and SO2 over FZN nanocatalyst. 3.5. Temperature programmed desorption of ammonia The results of NH3–TPD analyses for both the catalysts are shown in Fig. 3b. Two distinctive peaks observed in the NH3–TPD patterns of both the catalyst. The low temperature peak corresponds to the ammonia desorbed from weak acid sites at 140 °C, whereas the high temperature peak represents the ammonia desorbed from the strong acid sites in the temperature range of 200–350 °C [49]. In general it is opined that the weak acids sites may not contribute to the activity of the catalyst. It indicates that FZN nanocatalyst contains more broad range of strong acid sites compared to the FZ catalyst which is in line with the thermal desorption of DMS which may contribute to the activity of the catalysts [7,50]. The low-temperature desorption of DMS, in addition to high amount of strong acid sites are responsible for the high activity for oxidation of DMS with ozone at low temperature (<150 °C). 3.6. SEM and TEM analysis The surface morphology of FZ and FZN catalysts obtained by SEM are depicted in Fig. 4a and b, respectively. The Fe distribution is quite uniform on surface and pores morphology is lightly rough in case of FZ catalyst. FZ catalyst shows the particles aggregated into compact irregular shapes. At the same magnification, SEM image of FZN catalyst shows the particles are relatively less compact and smaller lumps can be observed. The Fe wt% obtained from EDAX for both the catalyst is presented in Table 1. The Transmission electron micrograms of FZ and FZN catalyst are presented in Fig. 4c and d, respectively. The size of the particle prepared by different methods in this study differs strongly, like co-precipitation method gave wider particle distributions. The Fe2O3 particles are observed by TEM analysis of FZN catalyst area more homogeneous and well-dispersed Fe2O3 distribution compare to poor Fe2O3 distribution in case co-precipitated FZ catalyst and some iron agglomeration are observed. It is found that the formed Fe2O3 nanostructures are more uniformly dispersed over the ZrO2 for FZN catalyst (Fig. 4d), and Fe2O3 nanoparticles with narrow size distribution are also observed. During the citrate synthesis process, an efficient interaction between citric acid and Fe2+ occurred, resulting in improved metal dispersion. The particle sizes of Fe are strongly depending on the citric acid and metal component interaction which prevent the aggregation of Fe nanoparticles. It would be beneficial for the formation of uniformly dispersed Fe nanoparticles on the support [48]. Though, various synthetic methods are reported for Fe2O3 nanoparticles, but wet chemical process

TDS (DMS) area 150 °C

380 °C

841 656

301 964

TPR area H2 uptake (cc)

TEM (nm)

8.5 10.2

120 17

is more efficient for producing the small size nanoparticles [31,32]. The size distribution of Fe2O3 particles for both the catalysts is depicted in Fig. 5a and b. For the catalysts synthesized by modified sol–gel method majority of the particles (40% of the particles) are in the range of 16–20 nm whereas in co-precipitated catalyst, they are in the range of 100–130 nm (60% of the particles). It is reported that during the thermal pre-treatment at 450 °C the carbon of the citrate is incompletely removed during calcinations process [36]. The thermal pre-treatment therefore leads to a presumably porous layer of intimately mixed iron oxide and carbon of a uniform thickness on the ZrO2 support. The carbon remaining after the thermal pre-treatment effectively prevents the formation of larger crystalline iron oxide particles. The nanoparticles obtained in this study are comparable to reported for 6 wt.% Fe/SiO2 catalyst prepared by colloidal method [48]. 4. Catalytic activity studies 4.1. Activity test for ozone decomposition Ozone decomposition experiments are performed under comparable conditions to ascertain the ozone decomposition on catalytic activity. The results are presented in Fig. 6a. The blank test (in the absence of any catalyst) results indicate that gaseous O3 decomposition is varied from 18% to 28% with increasing temperature from 50 to 200 °C; these results are in agreement with the observations of Shekar et al. [7]. However, in the presence of catalyst, the decomposition of ozone is increased significantly even at a relatively low temperature. With increase in temperature the Fe assisted ozone decomposition is significant due to the increase in surface potential energy of the catalyst. There are several lines of evidences that catalytic decomposition of ozone occurs on supported Fe2O3 and MnO2–Fe2O3 [8,51,52]. The ozone decomposition studies reveal that relatively higher amount of ozone decomposition is observed on FZN nanocatalyst compared to the co-precipitate FZ catalyst from 50 to 150 °C. In contrast, at 200 °C irrespective of the particle size both nano and co-precipitate catalysts performance is similar. The thermal effects are predominant at approaching 200 °C than particle size. It may be attributed to the higher surface area and better dispersion of active sites on FZN nanocatalyst compared to the FZ catalyst. On the other hand, the TPD of DMS and ammonia values for iron nanocatalyst exhibited stronger acidity as well as acid site population compared to that of co-precipitate catalyst. 4.2. Effect of temperature on DMS oxidation The effect of temperature on DMS oxidation is performed at ozone to DMS ratio of 1 and GHSV of 15,000 h 1. The DMS conversion as function of reaction temperature is displayed in Fig. 6b. Using oxygen as an oxidant, the DMS conversion values at 150 °C are attained a maximum value of 56% and 36% for FZN nanocatalyst and FZ catalyst, respectively. Further increase in reaction temperature led to little decline in conversion due to the sulfur poisoning during the oxidation of DMS over catalyst [53,54]. Several researchers have also been reported that the adsorption and activation of the O2 molecules over the catalyst surface require higher

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Fig. 4. SEM micrograph of (a) FZ and (b) FZN catalysts, TEM micrograph of (c) FZ and (d) FZN catalysts.

Fig. 5. Histogram of the particle size distributions as obtained from the TEM, (a) FZ and (b) FZN catalysts.

temperatures for catalytic incineration of methanethiol and DMS as described by the Langmuir–Hinshelwood model [4,55]. The oxidation of DMS with ozone is carried out in the absence of a catalyst also. The conversion value of DMS is about 23% at 150 °C. In the case of ozone and FZ catalyst, the minimal activity observed to be 72% at 50 °C and goes through a maximum value of 89% at 150 °C. However, it is interesting to observe that when ozone and FZN nanocatalyst both are employed in oxidation reaction, the conversion value of DMS is about 95% at 50 °C, and reaches maximum of 100% at 150 °C. Further increase in reaction temperature toward 200 °C slight decreases the conversion of DMS. The decline in the oxidation activity at higher temperature is attributed to the thermal decomposition of ozone which leads to the reduction in the

Fig. 6. Effect of reaction temperature (a) ozone decomposition and (b) conversion of DMS over FZ and FZN catalysts (GHSV = 15,000 h 1 and O3/DMS = 1).

amount of ozone available for catalytic decomposition [7,14]. The ozone decomposition studies are also indicating that thermal effects are predominant than particle size when approaching 200 °C. Catalytic oxidation with ozone offers new low-temperature reaction a way to achieve higher rates and lower activation energies than catalytic oxidation with molecular oxygen. In order to ascertain the synergetic effect of nanoparticles and ozone, the activation energy is calculated for both catalysts in ozone assisted catalytic oxidation of DMS. In the absence of ozone, the activation energies are 38 kJ mol 1 for FZN nanocatalyst and 44 kJ mol 1 for FZ catalyst, respectively. However, as ozone is added, the activation energy is significantly reduced to be 13.2 kJ mol 1 for FZN


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nanocatalyst which is smaller than the values obtained for FZ catalyst (22.8 kJ mol 1). Thus, the activation energy for the oxidation reaction is significantly reduced when instead of oxygen, ozone is applied. These values are similar to those reported previously for the oxidation of chlorobenzene over Iron oxide catalysts (20 kJ mol 1) [8]. Due to the smaller particles of iron oxide in FZN nanocatalyst, the activation energy for ozone assisted catalytic oxidation reaction is observed to be two folds lower than that of the for FZ catalyst which clearly demonstrates the advantage of nanoparticles. GC–MS analysis of non–condensable reaction samples at different reaction temperatures over FZN nanocatalyst are shown in Fig. 7. It is observed that SO2 and COx are the major products in the reactor effluent stream at higher temperature range from 100 °C to 200 °C. The effect of reaction temperature on product selectivity of DMS oxidation over FZN and FZ catalysts in presence of ozone are depicted in Fig. 8a and b, respectively. The FZ catalyst

shows that the selectivity to COx is observed minimum of 58% and goes through a maximum of 80% at 200 °C. The DMSO2/DMSO ratio increased as the reaction temperature increased, the DMSO fraction is reduced to less than 1% and the major partial oxidation product is DMSO2 at 150 °C. With further increase in temperature DMSO2 is decreased by 20% at 200 °C. In the presence of FZN catalyst, the COx selectivity is observed to be 80% and increased with reaction temperature and attained a stable value of 99% at 150 °C. On the other hand the selectivity to DMSO and DMSO2 are decreased with temperature and attained a stable value of 1%. In the co-precipitate catalyst the increase in the selectivity of DMSO2 could be due to the sulfur poisoning of active sites which are responsible for the conversion of DMSO2 to CO2 and SO2 though ozone is strong oxidizer [50]. In contrast, the stable activity of FZN nanocatalyst clearly indicates that the deactivation is completely minimized and nanoparticles size of Iron oxide offering the extended oxidation of DMSO2 to complete oxidation products such as CO2 and SO2. It is also observed that FZN catalyst contains more number of strong acid sites and higher H2 uptake capacity, which in turn reflects the number of active sites than FZ catalyst. The lowest temperature in TPR for Iron nanoparticles indicates the easy reducibility which means the faster re-oxidation of active site by ozone and increases the complete mineralization of DMS [50]. 4.3. Effect of ozone to DMS molar ratios on DMS oxidation

Fig. 7. GC–MS analysis of non-condensable reaction samples DMS oxidation at different reaction temperatures over FZN nanocatalyst (GHSV = 15,000 h 1 and O3/ DMS = 1).

In order to ascertain the effect of oxidant concentration on the DMS conversion for both the catalysts, a set of experiments were conducted by varying the ozone/DMS mole ratios from 0.2 to 1.4 at 150 °C. DMS conversion and selectivity to COx on FZN nanocatalyst are obtained about 81% at ozone/DMS ratio of 0. 2, and increasing ozone/DMS ratio to 1, almost complete conversion of DMS and only COx formation are observed (Fig. 9a). While in the presence of the FZ catalyst the COx selectivity (40%) is significantly lower than the DMS conversion (68%) at the same ratio (Fig. 9b) and the maximum conversion of DMS (95%) and COx selectivity (87%) values are observed at the ozone/DMS ratio of 1.4. At ozone/DMS ratio of 0. 2, about 56% of DMSO formation is observed on FZ catalyst. As the ozone/DMS ratio increased, DMSO fraction is reduced to less than 2% and the DMSO2 formation increased to 28% at a ratio of 1. Due to nanosized of FZN catalyst, the formation of DMSO and DMSO2 is completely minimized at ozone/DMS ratio

Fig. 8. Effect on product selectivity of DMS oxidation with ozone at different reaction temperature (a) FZN nanocatalyst (b) FZ catalyst (GHSV = 15,000 h 1 and O3/DMS = 1).

Fig. 9. Effect on conversion and product selectivity of DMS oxidation at various O3/ DMS molar ratios (a) FZN nanocatalyst (b) FZ catalyst (Reaction temp = 150 °C and GHSV = 15,000 h 1).

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of P1. The selectivity of the partial oxidation products is found to higher for FZ catalyst compared to the FZN nanocatalyst at lower ozone concentration. The nanocatalyst has larger contacting areas and resulted in more oxygen species formed during ozone decomposition and a higher conversion efficiency of DMS is also obtained. As shown by the TPR results, FZN nanocatalyst has a higher degree of reduction at low temperature, which may lead to higher lattice oxygen mobility and higher oxidation activity [56]. The total oxidation products are increased with higher ozone concentrations in alcohols such as ethanol and methanol over SiO2, and alumina [14,57]. It indicates that with increase in ozone concentration, the atomic oxygen species formed on the surface of the FZN nanocatalyst is relatively more and resulted in deeper oxidation for nanocatalyst in comparison to the FZ catalyst.

4.4. Effect of GHSV on DMS oxidation


4.5. Time on stream Times on stream analysis of DMS oxidation over both catalysts are studied at a reaction temperature of 150 °C, at Ozone/DMS ratio of 1 and GHSV of 15,000 h 1. The time as function of conversion and productivity are depicted in Fig. 11a and b, respectively. In this study it is observed that FZN nanocatalyst exhibits a total oxidation of DMS into COx and SO2 which exhibited constant activity for almost 3 h. FZN nanocatalyst shows that COx selectivity is almost 100% with time and no deactivation of catalyst are observed with long run time. The FZ catalyst activity is about 93% in the initial period and attained a steady state conversion value 71% in 3 h and the selectivity toward COX decreased from 78% to 55% in this period. In contrast to the FZ catalyst, the POP formation is also at minimal or negligible quantities on FZN catalyst. The product selectivity indicates that a fraction of the active sites are deactivated due to the sulfur poisoning and unable to convert the ozone in active species on FZ catalyst. A similar, rapid decline of the performance of Pt/Al2O3 catalyst has been observed due to the oxidation of DMS, forming irreversible sulfur-poisoning sites [54]. It can also be seen that the selectivity of partial oxidation product increased with extended reaction time for FZ catalyst. The smaller particles of iron oxide are enhancing the ozone decomposition as well as rate of DMS adsorption and decomposition. This study clearly demonstrates that the deactivation is completely minimized and synergetic effect of ozone and nanoparticles of FexOy are maintaining the surface active and the steady-state is achieved much faster compared to that of co-precipitate catalyst.

The effect of GHSV on DMS oxidation is studied at Ozone/DMS ratio of 1 and at 150 °C. The DMS conversion and product selectivity catalysts results are shown in Fig. 10a and b, respectively. The results reveal that the DMS conversion and COx selectivity values are decreased with increase in GHSV from 12,000 h 1 to 21,000 h 1. However, the decrease is less predominant in case of FZN nanocatalyst than FZ catalyst. In the presence of FZN nanocatalyst, the conversion of DMS and COx selectivity are almost constant 100% in the range of GHSV of 12,000–15,000 h 1. Further increase in GHSV up to 21,000 h 1, DMS conversion and COx selectivity are decreased to 90% and 87%, respectively. In this study, no DMSO2 formation is observed, only DMSO formation is observed over FZN nanocatalyst (<14%) at higher GSHV. In case of FZ catalyst, the conversion of DMS is decrease from 100% to 80% with increase in GHSV from 12,000 to 15,000 h 1. The selectivity of COx is observed to be 78% at GHSV of 15,000 h 1 and DMSO2 selectivity is about 18% at the same conditions. The ratio of DMSO/DMSO2 is increased at higher GHSV. Further increase in GHSV, the selectivity of DMSO and COx over FZ catalyst is observed around 50% at GHSV of 21,000 h 1. It demonstrates that the consecutive reaction to form COx is suppressed at higher GHSV and the overall yields of DMSO and COx remained almost constant which is well predicted in case of FZ catalyst. The increase in POP selectivity can be attributed to a decrease in contact time between reactant and the catalyst with increasing GHSV.

The products identified by online GC–MS analysis during catalytic oxidation of DMS with ozone are DMSO, DMSO2 carbon oxide and sulfur dioxide. Overall the carbon oxide and sulfur dioxide observed as major products of vapor phase catalytic ozonation. Though the literature reports [1,50] reveals that the formations of formaldehyde, methanethiol and methyl methanesulfonate are also observed but, these compounds may not stable in present conditions. Hence, these could not be detected in the present study due to the high ozone concentration and oxidizing power of iron catalyst. Detailed gas phase and heterogeneous oxidation of DMS

Fig. 10. Effects of conversion and product selectivity on DMS oxidation at various GHSV (a) FZN nanocatalyst (b) FZ catalyst (Reaction temp = 150 °C and O3/DMS ratio = 1).

Fig. 11. Effects of time on steam on conversion (Reaction temp = 150 °C, GHSV = 15,000 h 1 and O3/DMS ratio = 1).

5. Probable reaction pathway for DMS oxidation in presence of ozone


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under different condition are also reported [58,59]. A plausible reaction mechanism based on reaction products are detected in catalytic ozone reaction pathways of DMS degradation and mineralization (Scheme 1). It is also opined that ozone significantly alters the catalytic mechanisms which is fairly different from molecular oxygen. Mostly O3 assisted gas phase reactions involves radical chain mechanism initiated by O3 decomposition in presence of third body. The gas ozone does not completely oxidize organic compounds because low molecular weight partial oxidation products such as acetic acid and oxalic acid are relatively unreactive [8]. Moreover, in the gas phase, the observed reactions are likely due to the radical chain reactions initiated by the decomposition of ozone to O2 and O radical. It is proposed that not only ozone but also molecular oxygen may be involved in the autoxidation processes, in which the radical intermediates formed in VOC oxidation are oxidized by O2 [60]. However, when the concentration of the organic compounds is low, the chains do not propagate. The use of a catalyst offers an alternative pathway for ozone decomposition through a surface mechanism. This result has highlighted an interesting characteristic: a new active oxygen species may be formed on the surface of the catalysts during ozone decomposition and which is likely the dominant species for DMS oxidation at low temperatures. Oyama [9] have proposed that the adsorbed oxygen species are formed on the surface of MnOx catalysts during O3 decomposition. Adsorbed oxygen species formed from ozone decomposition on metal oxides is the main oxidizing species to convert VOCs due to its high oxidizing potential [61]. In situ Raman spectroscopy identified a band at 890 cm 1, which is an adsorbed peroxide species (O22 ) formed on the manganese oxide from ozone decomposition [62]. The previous results indicate that decomposition of ozone on a-Fe2O3 leads to the formation of highly reactive atomic oxygen species on the surface of the catalyst [63]. Konova et al. [64] proposed the mechanism of catalytic ozone decomposition on the transition metal oxides reported that the active complex of O [Co4+] are formed during the reaction of ozone decomposition on CoOx/Al2O3, which are capable of oxidizing VOCs at room temperature. Imamura et al. [60] has suggested on the basis of ESR studies that the oxygen anions O are formed on Ag2O and CeO2 in the presence of ozone. It can be seen that addition of ozone significantly increases the oxidation activity of FexOy catalysts and reduces the oxidation temperature of DMS (<200 °C) compared with gas phase decomposition alone. It may be due the free-radical species derived from thermal ozone decomposition react with DMS homogeneously as well as active oxygen species formed during the decomposition of ozone on FexOy catalyst and further react with

gaseous DMS molecules. It is suggested that highly reactive atomic oxygen species (O ) adsorbed play an important role of DMS oxidation. Based on the mechanisms proposed by Naydenov et al. [61], it is suggested that an active complex (O Fe3+) might be formed during ozone decomposition and which are capable of oxidizing DMS. Both homogeneous and heterogeneous reaction mechanisms may be present at the same time in the temperature range of 100– 200 °C. In present study the reactor is designed in such a way that the gas phase reaction is minimized. Previous study indicates that only at low temperature (<200 °C), sufficient ozone could reach the surface of the catalyst on which it could be decomposed, leading to formation of active oxygen species [65]. In this study, the ozone inlet position is close to the catalyst so that most ozone adsorbed on the catalyst decomposed thermally and reactive atomic oxygen would be formed from ozone decomposition on the catalyst even at elevated temperature. The decomposition of ozone is closely associated with the oxidation of DMS because both occur simultaneously on the surface. An analysis of the product confirmed an agreement with known information about the critical steps of ozone oxidation and the radical chemistry of organosulfur compounds. Literature reports reveal that there is a formation for methanesulfinic acid [CH3SO2H, MSIA], methanesulfonic acid [CH3SO3H, MSA] and sulfate as reaction products also observed during DMS oxidation [1,14,50]. Ozone is decomposed into oxygen species as a result of interactions with Fe3+, forming catalytic active sites O Fe3+. The initial reaction of DMS with surface active site O Fe3+ forms the DMSO, which on further oxidation generates DMSO2. The DMSO and DMSO2 formed will be oxidized further to MSIA and MSA intermediates, respectively. The ozone will rupture the –C–S– bond of these intermediates and further oxidation produces SO2 and COx. Under employed conditions MSIA and MSA are not observed which is one of the key steps in the formation of carbon oxides and SO2 which might be ascribed to the amount of reactive active O Fe3+ species available on the Fe–ZrO2 catalyst surface in addition to the reaction temperature. These intermediates appeared during the catalytic oxidation of DMS are also reported extensively in oxygen-based oxidation operations [58,59,66–70]. DMS reacts with OH, NO3, Cl and halogen oxide radicals in the atmosphere, providing SO2, DMSO, DMSO2, MSIA and MSA [71]. The present study reveals that the oxidation of DMS in which, cleavage of –C–S– bond is favorable than that of oxidation of –S– moiety at low reaction temperature in the presence of ozone. It may be attributed that ozone may act as a Lewis base, reacting with Lewis acid sites of FexOy as well as ZrO2 (unoccupied atomic orbital surface Fe or Zr atom). The FexOy and ZrO2 are both p-type oxides that exhibits abundant oxygen vacancies on its surface and

Scheme 1. Probable scheme of the reaction mechanism for the catalytic oxidation of DMS in presence of ozone.

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higher affinity to ozone [9,72]. It is also reported that the Fe3+ species produced by ion-exchange are stabilized on the surface of ZrO2 by forming Fe3+–O–Zr4+ structures [29]. At higher ozone concentrations the available surface adsorbed oxygen species is more and extended reactions proceeding to breaking of carbon sulfur bond. More than stoichiometric ratios of ozone and DMSO provided the complete mineralization which involves series of gas phase as well surface reactions and leading to the total oxidation. TPR results indicated that redox cycle of Fe catalysts require high temperature, it appears that ozone has been directly influencing the redox cycle at mild temperature. The physisorption of gaseous DMS on surface active site of O Fe3+ as well as zirconia is the first step which facilitates the cleavage –C–S– rather oxidation of –S– moiety and forms surfaced adsorbed intermediate species [–CH3O] and [–OSCH3]. Further oxidation of these surface adsorbed intermediates produces SO2 and COx. In this process catalyst surface Fe3+ is reduced to Fe2+. In the final step, gaseous ozone interact with Fe2+ as well as ziconia sites formed an active complex (O Fe3+) during ozone decomposition and which are capable of oxidizing DMS. It is also reported that the O initiated oxidation decomposed DMS and produced intermediate species such as [H3SCH2OO ], and then further converted them to stable products [67,71]. It is also reported that during the oxidation of DMS several intermediates such as, [CH3SCH2OO], [CH3S], [CH3SO] and [CH3SO2] are formed and further oxidation with ozone produces SO2 and COx [66–70]. It may be due to the fact that large amounts of active oxygen species formed during ozone decomposition which effectively oxidizes intermediates on the surface of the catalysts. It is observed that the selectivity of catalysts is governed by the nature of oxygen species; electrophilic oxygen species such as O2 and O are responsible for total oxidation of hydrocarbons, whereas the nucleophilic lattice stabilized O2 is responsible for the formation of partial oxidation products [73]. According to these analytical results, it is assumed that a small amount of partially oxidized byproducts could be formed and deposited on the surface of the iron oxide catalyst during DMS oxidation. It is reported that Fe/MgO nano-crystal catalysts are deactivated by transforming into the iron sulfide and are not reactive with the VOSC [74]. Fe-based nanocatalyst in a heterogeneous catalytic system have potential in the oxidation and decomposition of DMS because it can be regenerated by contacting with ozone. Furthermore, ozone is a powerful oxidant and may oxidize the intermediates and reduce surface poisoning [13,75].

6. Conclusions The iron nanocatalyst is successfully synthesized by modified sol–gel method and compared with the co-precipitate catalyst. The higher surface area and narrow pore size distribution are observed in nanocatalyst prepared by modified sol gel method. The XRD pattern revealed the amorphous nature of nanocatalyst and crystalline nature in co-precipitate catalyst. The lowest temperature in TPR for Iron nanoparticles indicates the easy reducibility which means the faster re-oxidation of active site by ozone. The TPD profile of DMS confirms the significant variation in the semi quantitative analysis of the active sites and acid sites observed. The TPD of ammonia results indicates that FZN nanocatalyst contains more broad range of strong acid sites compared to the FZ catalyst. TEM results indicate that average particle size of citrate synthesized iron oxide catalyst become much smaller (18 nm) compared with co-precipitated samples (120 nm). Higher amount of ozone decomposition is observed up to 150 °C for iron nanocatalyst compared to the co-precipitate catalyst, in contrast, at 200 °C irrespective of the particle sizes both nano and co-precipitate catalyst performance is similar. The thermal effects are


predominant at approaching 200 °C than particle size. The smaller particles of iron oxide are enhancing the ozone decomposition as well as rate of DMS adsorption and decomposition. This study clearly demonstrates that the deactivation is completely minimized and synergetic effect of ozone and nanoparticles of Fe are maintaining the surface active and the steady-state is achieved much faster compared to that of co-precipitate catalyst. It clearly demonstrates that under ozone assisted catalytic oxidation over nanocatalyst offering complete mineralization at 150 °C. The higher activity of FexOy nanoparticles in DMS oxidation is attributed to a small particle size, high surface area, high concentration of acid sites, and more densely populated surface coordination unsaturated sites. Acknowledgments I would like to thank Dr. K.S. Rama Rao, IICT, Hyderabad, Dr. G.K. Prasad, Scientist, DRDE, Gwalior and Prof. Sapna Sharma, JECRC Univeristy for providing valuable suggestions. The authors are grateful to the JECRC University, Jaipur, and Defence Research and Development Organization, Delhi, India for the partial support of this study. Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.jcis.2015.01.031. References [1] K. Demeestere, J. Dewulf, B.D. Witte, H.V. Langenhove, Appl. Catal., B 60 (2005) 93–106. [2] H. Li, J. Mihelcic, J. Crittenden, K. Anderson, J. Environ. Eng. 129 (2003) 684– 692. [3] S.W. Benson, Chem. Rev. 78 (1978) 23–35. [4] H. Chu, K. Horng, Sci. Total Environ. 209 (1998) 149–156. [5] H. Chu, W.T. Lee, Sci. Total Environ. 209 (1998) 217–224. [6] J.R. Kastner, K.C. Das, N.D. Melear, J. Hazard. Mater. 95 (2002) 81–90. [7] S.C. Shekar, K. Soni, R. Bunkar, M. Sharma, B. Singh, A. Nigam, T. Mahato, R. Vijayaraghavan, Catal. Commun. 11 (2009) 77–81. [8] H.C. Wang, H.S. Liang, M.B. Chang, J. Hazard. Mater. 186 (2011) 1781–1787. [9] S.T. Oyama, Catal. Rev. Sci. Eng. 42 (2000) 279–322. [10] C.B. Almquist, E.S. Demessie, K.S. Sehker, J. Sowash, Environ. Sci. Technol. 41 (2007) 4754–4760. [11] Y. Xi, C. Reed, Y.K. Lee, S.T. Oyama, J. Phys. Chem. B 109 (2005) 17587–17596. [12] A.R. Shawwa, D.W. Smith, Ozone Sci. Eng. 23 (2001) 161–170. [13] H. Einaga, S. Futamura, J. Catal. 227 (2004) 304–312. [14] E.S. Demessie, V.G. Devulapelli, Appl. Catal., B 84 (2008) 408–419. [15] E.S. Demessie, V.G. Devulapelli, Appl. Catal., A 361 (2009) 72–80. [16] M. Iwasaki, M. Hara, S. Ito, J. Mater. Sci. Lett. 17 (1998) 1769–1771. [17] P.V. Kamat, D. Meisel, Curr. Opin. Colloid Interface Sci. 7 (2002) 282–287. [18] D. Huang, D. Cao, Y. Li, H. Jiao, J. Phys. Chem. B 110 (2006) 13920–13925. [19] O. Shekhah, W. Ranke, A. Schule, G. Kolios, R. Schlogl, Angew. Chem. Int. Ed. Engl. 42 (2003) 5760–5763. [20] A.N. Pour, S. Taghipoor, M. Shekarriz, S.M.K. Shahri, Y.J. Zamani, Nanosci. Nanotech. 9 (2009) 4425–4429. [21] A.N. Pour, M.R. Housaindokht, S.F. Tayyari, J. Zarkesh, J. Nat. Gas Chem. 19 (2010) 284–292. [22] S. Eriksson, U. Nylen, S. Rojas, M. Boutonnet, Appl. Catal., A 265 (2004) 207– 219. [23] L. Chen, K. Zhu, L. Bi, A. Suchopar, M. Reicke, G. Mathys, H. Jaensch, U. Kortz, R.M. Richards, Inorg. Chem. 46 (2007) 8457–8459. [24] D. Habibi, A.R. Faraji, M. Arshadi, J.L.G. Fierro, J. Mol. Catal. A: Chem. 372 (2013) 90–99. [25] R. Srinivasan, T.R. Watkins, C.R. Hubbard, B. Davis, Chem. Mater. 7 (1995) 725– 730. [26] M. Waqif, J. Bachelier, O. Saur, J.C. Lavalley, J. Mol. Catal. 72 (1992) 127–138. [27] C. Morterra, E. Giamello, G. Cerrato, G. Centi, P. Perathoner, J. Catal. 179 (1998) 111–128. [28] L.A. Boot, A.J.V. Dillen, J.W. Geus, F.R.V. Buren, J. Catal. 163 (1996) 186–194. [29] Y. Okamoto, T. Kubota, Y. Ohto, S. Nasu, J. Catal. 192 (2000) 412–422. [30] J.M. Liu, P.Y. Lu, W.K. Weng, Mater. Sci. Eng., B 85 (2001) 209–211. [31] C. Solans, P. Izquierdo, J. Nolla, N. Azemar, M.J.G. Celma, Curr. Opin. Colloid Interface Sci. 10 (2005) 102–110. [32] L. Lu, K.H. Zhong, M. Wenling, L. Huajie, W. Yuqiu, J. Phys. Chem. 110 (2006) 15218–15223. [33] T. Sugimoto, K. Sakata, J. Colloid Interface Sci. 152 (1992) 587–590.


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