Microwave Assisted Synthesis of Nanostructured Titanium Dioxide with HighPhotocatalytic Activity

Anirudha Jena,† R. Vinu,‡ S. A. Shivashankar,† and Giridhar Madras*,‡

Materials Research Centre and Department of Chemical Engineering, Indian Institute of Science,Bangalore-560012, India

TiO2 (anatase) was synthesized using a microwave-irradiation-assisted chemical method. The reaction conditionswere varied to obtain unique nanostructures of TiO2 comprising nanometric spheres giving the materials avery porous morphology. The oxide was characterized by X-ray diffraction (XRD), Raman spectroscopy,transmission electron microscopy (TEM), scanning electron microscopy (SEM), X-ray photoelectronspectroscopy (XPS), Fourier transform infrared (FTIR) spectroscopy, and thermogravimetric analysis (TGA).The specific surface area and porosity were quantified by the BET method, and the degradation of dyes wascarried out using these materials. The photocatalytic activity of the nanometric TiO2 was significantly higherthan that of commercially available TiO2 (Degussa P25) for the degradation of the dyes.

Introduction

Nanostructured materials are of growing importance becauseof their many interesting properties and applications in differentfields. Nanosized titanium dioxide (TiO2) has been used for theenvironmental decontamination of a wide variety of organicdyes, bacteria, viruses, fungi, and so on.1,2 Although severalsemiconducting oxides have been used for such applications,TiO2 is the best choice as a photocatalyst because of its non-toxicity, low cost, and redox efficiency. The efficient photo-catalytic activity of TiO2 can be attributed to its electronic bandstructure. TiO2 is a large-band-gap material, and the band gapchanges upon transformation from the anatase (3.2 eV) to rutile(3.02 eV) to brookite (2.96 eV) form. This wide band gap ofTiO2 is responsible for its activity as a photocatalyst in thephotodecomposition organic pollutants.3,4

Savage et al. reported that TiO2 can also be employed formonitoring ambient gases, such as carbon monoxide, oxygen,and methane.5-7 A significant amount of work has been doneto study the activity of the titania catalyst, which depends onits crystal structure, morphology, specific surface area, andporosity, as well as on the shape and size distribution of pores.8,9

Because of such potentially wide-ranging applications, variousmethods have been employed to prepare nanostructures of TiO2

in powder as well as thin-film forms, as discussed in a recentreview.10 Anodic oxidation of Ti foil to obtain nanotubes ofTiO2 is one of the most studied processes.11,12 Other chemicalmethods to synthesize titania nanostructures include thesolvothermal,13,14 sol-gel,15 sol-gel-mediated dip-coating,16,17

reverse-micelle,18 microwave-assisted hydrothermal,19,20 mi-crowave-sintering,21 and microwave-assisted esterification22

processes.Although the formation of titania spheres has been reported

in the literature,23 the synthesis processes involve the use oftemplates,24 require a longer duration of reaction, or involveeither special conditions or sophisticated instruments such asspecially designed autoclaves25 that are capable of synthesizingmaterials under controlled pressure and temperature. Thus,obtaining uniformly distributed particles with a controlled phase

using a single-step surface modification in a simple microwaveoven is still a challenge.

In the present work, we describe the synthesis of well-crystallized TiO2 in the anatase phase in a simple microwaveoven, with a rather narrow size distribution in the nanometerrange. The spherical morphology of the titania samples is unique,especially because no templating or structure-directing agentwas used in the synthesis. The incorporation of hydroxyl groupson the surface of the oxide material was achieved through anovel synthesis procedure. The photocatalytic degradation ofseveral anionic and cationic dyes using the synthesized materialwas studied and compared with the degradation obtained in thepresence of commercially available titania (Degussa P25,denoted DP25).

Experimental Procedure

Analytical-reagent- (AR-) grade chemicals were used in thesynthesis. The chemicals were obtained from various sources:Ti(iOPr)4 (iOPr ) isopropopoxide, 97%) from Sigma-Aldrich(St. Louis, MO); the surfactant poly(vinylpyrrolidone) (PVPK-90) from Rolex (Mumbai, India); the dyes Orange G (OG),Remazol Brill Blue (RBB), Alizarin Red S (ARS), MethyleneBlue (MB), and Rhodamine B (RB) from S.D. Fine ChemicalsLtd. (Mumbai, India). HPLC-grade ethanol (99.9%) was ob-tained from Les Alcools de Commerce (Boucherville, QC,Canada) and was further dried using the standard procedure.26

Hydrazine hydrate was obtained from Rankem (New Delhi,India).

Synthesis of TiO2. Nanocrystallites of titania powder weresynthesized by a microwave-irradiation-assisted chemical pro-cess of the type reported recently.27-29 In the present work onthe synthesis of TiO2 from Ti(iOPr)4, we used two differentconditions. In the first batch, 2.94 mL of Ti(iOPr)4 was dissolvedin 30 mL of dried ethanol (dielectric constant, ε′ ) 24.3), andin the second batch, 5 mL of hydrazine hydrate (HH, ε′ ) 51.7)was added to a solution of Ti(iOPr)4 in ethanol, prepared in thesame manner as in the first batch. The oxide samples obtainedfrom these two batches are referred to as TiO2(E) and TiO2(HH),respectively. To each of these solutions was added a solutionof the surfactant poly(vinyl pyrrolidone) (PVP, molecular weight) 360000; 0.3 g of PVP dissolved in 15 mL of ethanol), andthe consolidated (reactant) solution was stirred for 30 min. The

* To whom correspondence should be addressed. E-mail: giridhar@chemeng.iisc.ernet.in. Tel.: +91 80 2293 2321. Fax: +91 80 2360 0683.

† Materials Research Centre.‡ Department of Chemical Engineering.

Ind. Eng. Chem. Res. 2010, 49, 9636–96439636

10.1021/ie101226b 2010 American Chemical SocietyPublished on Web 09/28/2010

reactant solution, taken in a round-bottomed flask, was subjectedto microwave irradiation at a power of 800 W in a domesticmicrowave oven (2.45 GHz) equipped with a reflux condenserplaced outside the oven.29 The duration of irradiation was 5min, following which the resulting colloidal solution was cen-trifuged to collect the solid powder product, which was heatedin air at 500 °C for 1 h to remove the surfactant and otherorganic byproducts of the irradiation-induced chemical reac-tion(s). The white powder so obtained was subjected to furtherinvestigation.

Photocatalytic Measurements. Photocatalytic measurementswere carried out in a reactor consisting of a cylindrical jacketedborosilicate glass container [5.6-cm internal diameter (i.d.), 7.6-cm outer diameter (o.d.), and 10-cm height] in which the dyesolution was degraded. UV radiation was provided by a 125-Whigh-pressure mercury vapor lamp placed inside a jacketedquartz tube of 4-cm i.d., 4.7-cm o.d., and 18-cm height. Thequartz tube was dipped into the solution, and the solution wasstirred continuously using a magnetic stirrer for uniform mixing.Cold water was circulated in the jacket of the container and inthe annulus of the quartz tube to quench the heat generated andto maintain the temperature at 35 °C. The photon flux of thesystem was 30 W/cm2. Further details of the experimental setupcan be found elsewhere.30

The degradation was carried out using an initial concentrationof 50 ppm for all five dyes, OG, RBB, ARS, MB, and RB. Itwas first ensured that no degradation occurred in the presenceof UV light but in the absence of the catalyst to ensure theabsence of photolysis. Then, 100 mL of dye solution and 100mg of TiO2 were stirred for 2 h in the absence of UV light tocheck the initial surface adsorption. The solution was thenexposed to UV light, and samples were collected at regularintervals. These samples were then centrifuged to remove thecatalyst. The maximum absorption wavelengths (λmax) for OG,RBB, ARS, MB, and RB as calculated from absorbance mea-

surements were 484, 592, 540, 665, and 554 nm, respectively.The change in concentration with respect to time for each dyeafter degradation was quantified, based on the calibration ofeach dye by spectrophotometry (Shimadzu UV 1700 Phar-maSpec). All degradation experiments were repeated, and theerror in the concentration was less than 5%. Further, experimentswere conducted with the same sample multiple times to checkthe lifetime activity. There was no decrease in the reaction rateafter five repeated experiments, indicating that the synthesizedcatalyst was stable.

Characterization of TiO2. The crystallinity of the white solidpowder (resulting from microwave irradiation) was studied ona Bruker D8 Advance X-ray diffractometer using Cu KR

radiation. Raman studies were carried out using argon ion laserof wavelength 514 nm with power (0.9 mW) supplied by aHoriba Jobinyvon LabRAM HR 100 Raman instrument. Mor-phological analysis was carried out using a field-emissionscanning electron microscope (FESEM-SIRION XL-40) operat-ing between 200 V and 30 keV. High-resolution transmissionelectron microscopy (TEM) images were obtained in a TEMinstrument operating at 30 keV (TECNAI F-30). The specificsurface area of the oxide samples was measured by theBrunauer-Emmett-Teller (BET) method (Micromeritics ASAP2020 system). Core-level spectral analysis of the synthesizedtitania powder was carried out on an X-ray photoelectronspectroscopy (XPS) instrument with a resolution of 0.8 eV(Thermo Scientific Multilab-2000). Fourier transform infrared(FTIR) spectra of the samples in KBr pellets were obtained inthe spectral range of 4000-400 cm-1 at a resolution of 4 cm-1

(Perkin-Elmer Spectrum GX). The thermal stability of the oxideswas studied by thermogravimetric analysis (TA Instrumentsmodel STQ6000). The band gap of TiO2 was determined from

Figure 1. X-ray diffraction (XRD) patterns of TiO2(E) and TiO2(HH).

Figure 2. Raman spectra of TiO2(E) and TiO2(HH).

Figure 3. SEM images of (a) TiO2(E) and (b) TiO2(HH).

Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010 9637

diffuse-reflectance measurements (Perkin-Elmer Lambda 35spectrophotometer).

Results and Discussion

X-ray Diffraction Analysis. The powder X-ray diffractionpatterns of both the oxides TiO2(E) and TiO2(HH) are shownin Figure 1. The patterns demonstrate clearly that both of thesamples are well crystallized. The patterns could be indexed totitania of the anatase phase (JCPDS card number 04-0477). Theabsence of peaks at Bragg angles (2θ) of 27.5°, 39.3°, and 54.2°confirms the absence of the rutile phase in the samples.Furthermore, the patterns were analyzed for phase content usingTopas 3 (Bruker) software, from which it was concluded thatboth samples consisted purely of anatase titania.

Raman Studies. The Raman spectrum of both TiO2(E) andTiO2(HH) are shown in Figure 2. According to factor groupanalysis, anatase has six Raman-active modes (A1g + 2B1g +3Eg).

31 The six allowed modes appear at 144 cm-1 (Eg), 197cm-1 (Eg), 399 cm-1 (B1g), 513 cm-1 (A1g), 519 cm-1 (B1g),and 639 cm-1 (Eg). Raman studies again indicated the absenceof any rutile phase in both the samples.

Electron Microscopy. FESEM images of the two samplesof TiO2 are shown in Figure 3. Although they appear to bebroadly similar agglomerations of crystallites, the finer details

of the microstructures and the distinctive characteristics arerevealed by transmission electron microscopy (Figure 4). Themorphology of the sample prepared with ethanol as the solventis striking, as the sample is made of very spherical agglomera-tions of smaller crystallites (Figure 4a). The spheres vary indiameters from about 40 to about 100 nm, with each spheremade of crystallites measuring about 7-10 nm. These smallercrystallites are not necessarily spherical and appear to be quitedensely packed in each sphere, thereby reducing the porositywithin each spherical assembly. Such nearly spherical ag-glomerations were found everywhere in the sample and werefound to be entirely reproducible.

Although somewhat spherical agglomerations of smaller TiO2

crystallites can be spotted in the titania sample prepared withhydrazine hydrate added to ethanol as the solvent (Figure 4b),the morphology of this sample is quite different. As can be seenfrom the TEM image, the crystallites in this sample are ratherloosely packed. This makes the assembly more porous than theassembly of spheres formed when ethanol was used as the onlysolvent during synthesis. The interstitial voids in the porousmicrostructure of TiO2(HH) suggest that this structure mightbe favorable to the photocatalytic activity of the oxide. Theinsets of Figure 4a,b show the selected-area electron diffractionpatterns of the oxide samples. The rings in the SAED patterns

Figure 4. TEM images of (a) TiO2(E) and (b) TiO2(HH), with selected-area electron diffraction patterns shown in the insets. HRTEM images of (c) TiO2(E)and (d) TiO2(HH).

9638 Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010

confirm that the samples are well crystallized. The interplanard spacings as calculated from the SAED patterns are 3.50, 2.4,1.89, and 1.67 Å, confirming that the titania is of the anatasephase. The high crystalline nature of both titania samples wasagain confirmed from the HRTEM images, as shown in Figure4c,d. From the HRTEM images, the crystallographic spacingwas calculated for both samples. The distinct lattice fringes of3.5 and 2.4 Å corresponds to the (101) and (103) planes, againconfirming the anatase phase.

Surface Area. Nitrogen sorption at 77 K was used for thedetermination of the specific surface areas of the samplesaccording to the BET method. The method was also employedto quantify the porosities of the two titania samples. The specificsurface areas of TiO2(E) and TiO2(HH) were found to be 47and 62 m2/g, respectively. The Barrett-Joyner-Halenda (BJH)pore size distribution plots are shown in the Figure 5a. Althoughthe two oxides TiO2(E) and TiO2(HH) have comparable surfaceareas, there is a significant difference in the pore sizes of thetwo materials. The BJH plots are based on isotherms of typeIV for N2 adsorption-desorption with hysteresis loops (Fig-ure 5b), clearly indicating the mesoporous nature of the TiO2.The two oxides were found to be mesoporous with mean porediameter of 16 and 41 nm, respectively. As shown in thecorresponding TEM image (Figure 4a), the smaller 16-nm poresin the case of TiO2(E) are due to the relatively tightly packedparticles that form each sphere. In contrast, in the case ofTiO2(HH), the 41-nm pores can be attributed to the interstitialvoids resulting from the relatively loose packing of thenanocrystallites. Such mesoporous titanias have been found toexhibit high photocatalytic activity and hence can potentiallybe used in biomemitic photocatalysis and dye-sensitized solarcells.32,33

XPS Analysis. Core-level spectral analysis was carried outfor both TiO2(E) and TiO2(HH). The intensity distribution forO 1s, as shown in the Figure 6a, results in a considerabledifference between TiO2(E) and TiO2(HH). The spectral line israther more intense and broad in the O 1s spectrum of TiO2(HH)

than in that of TiO2(E). The O 1s spectrum was deconvolutedusing XPS peak41 software. Upon fitting, the spectra show adoublet with well-resolved peaks at 530.8 and 532.4 eV for bothTiO2(E) and TiO2(HH) (Figure 6b,c, respectively). The 530.8eV peak can be assigned to the O2- from TiO2, whereas thehigher-binding-energy peak at 532.4 eV can be assigned to theadsorbed hydroxyl group on the oxide surface.34,35 The differ-ence between the binding energies of the assigned hydroxyl(OH-) and oxide (O2-) species was found to be 1.6 eV, whichis close to reported differences36-38 of 1.5-1.9 eV. Thedeconvoluted spectrum shows that the area under the curve forthe hydroxyl peak in TiO2(HH) is greater than that in TiO2(E)(Figure 6d). This indicates that the hydroxyl content ofTiO2(HH) is higher than that of TiO2(E). Parts e and f of Figure6 show Ti 2p and C 1s core-level spectra for TiO2(E) andTiO2(HH), respectively.

Band Gap and IR Measurements. The band gaps of thetwo oxide materials were calculated using the Kubelka-Munktreatment of the diffuse-reflectance (DR) spectra. The theory issimple and straightforward, as mentioned in earlier reports.39

The Kubelka-Munk equation at any wavelength is F(Ri) ) (1- R)2/(2R), where F(Ri) is the Kubelka-Munk function and Ris the reflectance from the sample. The DR spectra of both thesamples are shown in Figure 7. A linear drop in the reflectanceis observed at a wavelength of around 380 nm, correspondingto an energy of 3.3 eV. Based on the Kubelka-Munk treatment(inset of Figure 7), it is concluded that the intersection betweenthe linear fit and the photon energy axis gives the value of Eg.The two materials show essentially identical band gaps, 3.32and 3.29 eV for TiO2(E) and TiO2(HH), respectively.

The FTIR spectra of the oxides TiO2(E), TiO2(HH), and DP25were collected in the frequency range of 4000-400 cm-1. Figure8 shows ν(OH) stretching bands in the region between 3700and 3300 cm-1. The bands observed at 3400 cm-1 are broadand are due to hydroxyl groups on different sites.34 The resultantspectral broadening could be from a superposition of thevibration bands of hydroxyl groups and the stretching vibrationsof adsorbed water molecules, which also can be confirmed fromthe band at 1636 cm-1 (i.e., due to the bending modes ofmolecular water).40 Clearly, there are noticeable differencesamong the spectra of TiO2(E), TiO2(HH), and DP25. There isa considerable difference between two oxide samples in the“area under the curve” for the very large band at 3400 cm-1

and the band at 1640 cm-1. In the case of TiO2(HH), thebroadening for the surface hydroxyl group and also the areaunder the OH band are highest, in agreement with the observedhigher catalytic activity.

Thermal Analysis of the Oxide. To examine the thermalstabilities and surface hydroxyl contents of the oxide samples,thermogravimetric analysis was carried out under flowing argon.The results for the two oxides are shown in Figure 9, whichindicates a higher weight loss for TiO2(HH). This can beattributed to the higher percentage of surface hydroxyl groupson the oxide surface. The hydrated form of hydrazine might bea cause of such incorporation of hydroxyl groups on the oxidesurface. The analysis shows about 7% weight loss for TiO2(HH),3% weight loss for TiO2(E), and 1.8% weight loss for DP25.

Microwave Synthesis Mechanism. Although the mechanismof synthesis of such unique spherical titania samples bymicrowave irradiation is not clear, based on the TEM micro-graphs, it can be suggested that the dielectric power and natureof the solvents determined the morphologies and porosities ofTiO2(E) and TiO2(HH). In the case of TiO2(E), after the titani-um isoproxide in ethanol is mixed with PVP, the polymer is

Figure 5. BJH pore size distributions of (a) TiO2(HH) and (b)TiO2(E). N2

adsorption and desorption isotherm.

Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010 9639

adsorbed onto the alkoxide surface, resulting in sphericalmicelles. When this mixture is further subjected to microwaveradiation, the initial clear solution becomes turbid, and finally,precipitation occurs, with the formation of titania. The initialtransparent solution indicates that no hydrolysis of the alkoxideoccurs before microwave irradiation. During the precipitationreaction, primary particles nucleate initially from the solution,

and these primary particles aggregate to form secondary par-ticles. The secondary agglomerates result in spherical aggregatesof nanosized crystalline titania.

In the case of TiO2(HH), the reaction mechanism is similarto that for hydrolysis.25 The titanium isopropoxide reacts withwater from the hydrated solvent, which further undergoes hy-drolysis and condensation reaction. Because the HH solvent is

Figure 6. XPS core-level analyses: (a) O 1s spectra of both TiO2(HH) and TiO2(E), (b) O 1s deconvoluted spectrum for TiO2(E), (c) O 1s deconvolutedspectrum for TiO2(HH), (d) -OH spectra for both TiO2(HH) and TiO2(E). Ti 2p and C 1s core-level spectra of (e) TiO2(E) and (f) TiO2(HH).

Figure 7. Diffuse-reflectance spectra of TiO2 for both TiO2(E) andTiO2(HH). Inset: Kubelka-Munk-transformed reflectance spectra of TiO2(E)and TiO2(HH). Figure 8. FTIR analyses of TiO2(E), TiO2(HH), and DP25.

9640 Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010

hydrated, the PVP chain is no longer linear, and the water fromthe solvent forms H-bonds with the PVP chain. This disturbsthe formation of spherical micelles, and therefore, agglomerationtends to organize the crystallites with voids and creates amesoporous structure.

The particle growth after the nucleation of the primaryparticles can be affected by the solvents, by the microwavepower, and also by the surfactant used during the microwave

treatment. The dielectric power of the solvent plays a crucialrole during the nucleation process. The dielectric constants ofethanol and HH are 24.3 and 51.7, respectively. These solventsinteract strongly with microwave power and form the sphericalaggregates of crystalline mesoporous nanotitania.25 Thus, thesolvent plays an important role in maintaining the sphericalshape of the particles and the mesoporous structure. Anotherreason for such a high degree of crystallization under microwaveheating must be bursting nucleation due to rapid and homoge-neous dielectric heating, which cannot be achieved by theconventional heating process.

Photocatalytic Degradation. The mechanisms for the pho-tocatalytic reactions of organic dyes have been well establishedby various authors.41-44 Because the absorption of light is afunction of concentration, the rate of degradation due tophotocatalysis of each dye can be calculated by conducting aseries of measurements on the change in concentration of thedye with respect to time. To enable direct comparison of thephotocatalytic performance of different samples of TiO2, thecatalyst loading was kept constant at 0.1 g per 100 mL ofsolution for all degradation measurements. The initial concentra-tion was chosen to be 50 ppm in all cases. However, adsorption

Figure 10. Photocatalytic degradations of (a) Orange G, (b) Remazol Brill Blue, (c) Alizarin Red S, (d) Methylene Blue, and (e) Rhodamine B in thepresence of three catalysts: TiO2(E), TiO2(HH), and DP25.

Figure 9. Thermogravimetric analyses of TiO2(E) and TiO2(HH).

Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010 9641

occurred during stirring in the dark for 2 h, and the initialconcentration decreased in certain cases. The concentrationobtained after stirring in the dark was therefore taken as C0.The degradation of the different dyes is represented graphicallyin Figure 10, which shows the changes in concentration of thedyes upon exposure to UV light in the presence of the titaniacatalyst. Indeed, a significant reduction in dye concentrationoccurred in less than 1 h of UV irradiation, with saturationreached at a low level of concentration in 1 h. In all of thedegradation studies with the microwave-synthesized titania, thefinal concentration of the dye reached below 10 ppm. From allof the results of the degradation experiments, it is clear that theTiO2 obtained by microwave irradiation has a higher activitythan the commercial catalyst DP25.

To further quantify the reaction, the reaction rate constantswere calculated. For a first-order degradation, the rate of thedegradation is given by -r ) dC/dt ) kC, where C is the dyeconcentration. For an initial rate analysis, the rate constant, k,can be determined from the above relation as k ) (C0 - Ct)/(tC0), where Ct is the concentration of dye after t ) 30 min ofphotocatalysis. The calculated rate constants for the varioussystems are listed in Table 1. From the observed rate constantvalues, it is clear that the degradation of the dyes follows theorder TiO2(HH) > TiO2(E) > TiO2 (DP25) for all of the dyesexcept RB. For the degradation of RB, the degradation rate ishigher for TiO2(E) than for TiO2(HH). However, regardless ofwhether a dye is cationic or anionic, the degradation of the dyewas found to be faster when the catalyst employed had beensynthesized by microwave irradiation.

The enhancement in the photocatalytic activity of variousmaterials relies mostly on the crystallinity, specific surface area,and pore size distribution. A high surface area offers more activeadsorption sites and leads to more photocatalytic reactioncenters, whereas porosity in a material helps in the formationof channels for the flow of reacting species.45 Chemical reactionsare effective when the transport paths through which moleculesmove are through these porous channels. The transport of smallmolecules in media featuring large mesopores can approachdiffusion rates comparable to those observed in open media.46

It is therefore believed that macrochannels in a sample can actas effective transport paths for reactants and also help toovercome intradiffusional resistance in a typical mesoporoustitania that consists of monolithic particles. The photoadsorptionefficiency is strongly influenced by the pore wall structure ofthe photocatalyst, and porous channels act as light-transfer pathsfor introducing incident photon flux onto the inner surface ofmesoporous TiO2 and improve the photoabsorption efficiencyof the catalyst.47 The interconnected TiO2 nanoparticle arraysembedded in the mesoporous wall allow highly efficientphotogenerated electron transport, photoadsorption, and transportof molecules through the macrochannel network. Thus, the useof a solvent (such as hydrazine hydrate) that both has a highdielectric constant and also facilitates hydroxyl group incorpora-tion on the surface yields TiO2 with a higher photocatalytic

efficiency. These features help explain the high photocatalyticactivity of the mesoporous titania samples obtained in this study.

Conclusions

Phase-pure, nanocrystalline anatase TiO2 was synthesized bymicrowave irradiation of an ethanolic solution of titaniumisopropoxide, with and without added hydrazine hydrate.Anatase of a unique microstructure, comprising densely packedspherical agglomeration of nanocrystalline TiO2, was obtainedwhen hydrazine hydrate was absent from the solution. Thepresence of HH led to a more mesoporous morphology, andthe porosity of the samples could be controlled by optimizingthe reaction conditions. The photocatalytic action of titania wasmeasured through the degradation of different organic dyes. Theuse of a hydrated and high-dielectric solvent such as hydrazinehydrate results in a material that shows better photocatalyticactivity than a commercial catalyst. Moreover, the observedspherical and mesoporous morphology, porosity, and catalyticactivity are entirely reproducible.

Acknowledgment

We thank the Institute Nanoscience Initiative for providingaccess to the electron microscopy facility, the Central X-rayfacility for XRD data, the Department of Inorganic and PhysicalChemistry for surface area measurements, the Centre forExcellence in Nanoelectronics for Raman measurements, theInstitute surface characterization facility for XPS data, and ProfS. Natarajan for diffuse-reflectance measurements. G.M. thanksthe department of science and technology for the Swarnajayanthifellowship.

Literature Cited

(1) Huang, Z.; Maness, P. C.; Blake, D. M.; Wolfrum, E. J.; Smolinski,S. L.; Jacoby, W. A. Bactericidal mode of titanium dioxide photocatalysis.J. Photochem. Photobiol. A: Chem. 2000, 130, 163.

(2) Senogles, P. J.; Scott, J A.; Shaw, G.; Stratton, H. Photocatalyticdegradation of the cyanotoxin cylindrospermopsin, using titanium dioxideand UV radiation. Water Res. 2001, 35, 1245.

(3) Patsoura, A.; Kondarides, D. I.; Verykios, X. E. Photocatalyticdegradation of organic pollutants with simultaneous production of hydrogen.Catal. Today 2007, 124, 94.

(4) Bahnemann, W.; Muneer, M.; Haque, M. M. Titanium dioxide-mediated photocatalysed degradation of few selected organic pollutants inaqueous suspensions. Catal. Today 2007, 124, 133.

(5) Savage, N.; Akbar, S. A.; Dutta, P. K. Titanium dioxide based hightemperature carbon monoxide selective sensor. Sens. Actuators B 2001, 72,239.

(6) Frank, M. L.; Fulkerson, M. D.; Patton, B. R.; Dutta, P. K. TiO2-based sensor arrays modeled with nonlinear regression analysis forsimultaneously determining CO and O2 concentrations at high temperatures.Sens. Actuators B 2002, 87, 471.

(7) Savage, N.; Chwieroth, B.; Ginwalla, A.; Patton, B. R.; Akbar, S. A.;Dutta, P. K. Composite n-p type semiconducting titanium dioxide as gassensors. Sens. Actuators B 2001, 79, 17.

(8) Chen, X.; Mao, S. Titanium dioxide nanomaterials: Synthesis,properties, modifications, and applications. Chem. ReV. 2007, 107, 2891.

(9) Thompson, T. L.; Yates, J. T., Jr. Surface science studies of thephotoactivation of TiO2sNew photochemical processes. Chem. ReV. 2006,106, 4428.

(10) Hu, X.; Li, G.; Yu, J. C. Design, Fabrication, and Modification ofNanostructured Semiconductor Materials for Environmental and EnergyApplications. Langmuir 2010, 26, 3031.

(11) John, S. E.; Mohapatra, S. K.; Misra, M. Double-wall anodic titaniananotube arrays for water photooxidation. Langmuir 2009, 25, 8240.

(12) Sreekantan, S.; Hazan, R.; Lockman, Z. Photoactivity of anatase-rutile TiO2 nanotubes formed by anodization method. Thin Solid Films 2009,518, 16.

Table 1. Rate Constants (× 10-3 min-1) for PhotocatalyticDegradation in the Presence of DP25, TiO2(E), and TiO2(HH)

dyes

catalyst OG RRB ARS MB RB

DP25 3.3 2.3 4.0 9.7 8.3TiO2(E) 10.6 4.6 8.9 16.8 14.0TiO2(HH) 20.4 5.2 15.0 18.4 11.6

9642 Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010

(13) Dinh, C. T.; Nguyen, T. D.; Kleitz, F.; Do, T. Shape-controlledsynthesis of highly crystalline titania nanocrystals. ACS Nano 2009, 3, 11–3737.

(14) Zhang, X.; Ge, X.; Wang, C. Synthesis of titania in ethanol/aceticacid mixture solvents: Phase and morphology variations. Cryst. Growth Des.2009, 9, 4301.

(15) Hussain, M.; Ceccarelli, R.; Marchisio, D.; Fino, D.; Russo, N.;Geobaldo, F. Synthesis, characterization, and photocatalytic application ofnovel TiO2 nanoparticles. Chem. Eng. J. 2010, 157, 45.

(16) Sun, P.; Liu, H.; Yang, H.; Fu, W.; Liu, S.; Li, M.; Sui, Y.; Zhang,Y.; Li, Y. Synthesis and characterization of TiO2 thin films coated on metalsubstrate. Appl. Surf. Sci. 2010, 256, 3170.

(17) Calabria, A. J.; Vasconcelos, W. L.; Daniel, D. J.; McPhail, D.;Boccaccini, A. R. Synthesis of sol-gel titania bactericide coatings on adobebrick. Constr. Build. Mater. 2010, 24, 384.

(18) Lim, K. T.; Hwang, H. S.; Ryoo, W.; Johnston, K. P. Synthesis ofTiO2 nanoparticles utilizing hydrated reverse micelles in CO. Langmuir2004, 20, 2466.

(19) Zhang, P.; Yin, S.; Sato, T. Synthesis of high-activity TiO2

photocatalyst via environmentally friendly and novel microwave assistedhydrothermal process. Appl. Catal. B 2009, 89, 118.

(20) Guobin, M. A.; Zhao, X.; Zhu, J. Microwave hydrothermal synthesisof rutile TiO2 nanorods. Int. J. Mod. Phys. B 2005, 19, 2763.

(21) Uchida, S.; Tomiha, M.; Masaki, N.; Miyazawa, A.; Takizawa, H.Preparation of TiO2 nanocrystalline electrode for dye-sensitized solar cellsby 28 GHz microwave irradiation. Sol. Energy Mater. Sol. Cells 2004, 81,135.

(22) Li, Y; Li, H; Li, T; Li, G; Cao, R. Facile synthesis of mesoporoustitanium dioxide nanocomposites with controllable phase compositions bymicrowave-assisted esterification. Microporous Mesoporous Mater. 2009,117, 444.

(23) Keshmiri, M.; Troczynski, T. Synthesis of narrow size distributionsub-micron TiO2 spheres. J. Non-Cryst. Solids 2002, 311, 89.

(24) Chen, J.; Hua, Z.; Yan, Y.; Zakhidov, A.; Baughman, R. H.; Xu,L. Template synthesis of ordered arrays of mesoporous titania spheres.Chem. Commun. 2010, 46, 1872.

(25) Periyat, P; Leyland, N; McCormack, D. E; Colreavy, J; Corr, D;Pillai, S. C. Rapid microwave synthesis of mesoporous TiO2 for electro-chromic displays. J. Mater. Chem. 2010, 20, 3650.

(26) Furniss, B. S.; Hannaford, A. J.; Smith, P. W. G.; Tatchell, A. R.Vogel’s Text Book of Practical Organic Chemistry, 5th ed.; Prentice Hall:Upper Saddle River, NJ, 1996.

(27) Bilecka, I.; Djerdj, I.; Niederberger, M. One-minute synthesis ofcrystalline binary and ternary metal oxide nanoparticles. Chem. Commun.2008, 7, 886.

(28) Brahma, S.; Shivashankar, S. A. Surfactant-mediated synthesis offunctional metal oxide nanostructures via microwave irradiation-assistedchemical synthesis. Mater. Res. Soc. Symp. Proc. 2009, 1174, 1174V02-08.

(29) Brahma, S.; Rao, K. J.; Shivashankar, S. A. Rapid growth ofnanotubes and nanorods of wurtzite ZnO through microwave-irradiation ofa metalorganic complex of zinc and a surfactant in solution. Bull. Mater.Sci. 2010, 33, 1.

(30) Mahata, P.; Madras, G.; Natarajan, S. Novel photocatalysts for thedecomposition of organic dyes based on metal-organic framework com-pounds. J. Phys. Chem. B 2006, 110, 13759.

(31) Choi, H. C; Jung, Y. M; Kim, S. B. Size effects in the Ramanspectra of TiO2 nanoparticles. Vibr. Spectrosc. 2005, 37, 33.

(32) Yu, J; Su, Y; Cheng, B. Template-Free Fabrication and EnhancedPhotocatalytic Activity of Hierarchical Macro-/Mesoporous Titania. AdV.Funct. Mater. 2007, 17, 1984.

(33) Kim, Y. J.; Lee, M. H.; Kim, H. J.; Lim, G.; Choi, Y. S.; Park,N. G.; Kim, K.; Lee, W. I. Formation of highly efficient dye-sensitizedsolar cells by hierarchical pore generation with nanoporous TiO2 spheres.AdV. Mater. 2009, 21, 3668.

(34) Erdem, B.; Hunsicker, R. A.; Simmons, G. W.; Sudol, E. D.;Dimonie, V. L.; El-Aasser, M. XPS and FTIR surface characterization ofTiO2 particles used in polymer encapsulation. Langmuir 2001, 17, 2664.

(35) Kurtz, R. L.; Stockbauer, R.; Madey, T. E. Synchrotron radiationstudies of H2O adsorption on TiO2(110). Surf. Sci. 1989, 218, 178.

(36) Simmons, G. W.; Beardt, B. C. Characterization of acid-baseproperties of the hydrated oxides on iron and titanium metal surfaces. J.Phys. Chem. 1987, 91, 1143.

(37) Sanjines, R.; Tang, H.; Berger, H.; Gozzo, F.; Margaritondo, G.;Levy, F. Electronic structure of anatase TiO2 oxide. J. Appl. Phys. 1994,75, 2945.

(38) Raikar, G. N.; Gregory, J. C.; Ong, J. L.; Lucas, L. C.; Lemons,J. E.; Kawahara, D.; Nakamura, M. Surface characterization of titaniumimplants. J. Vac. Sci. Technol. 1995, 13, 2633.

(39) Morales, A. E.; Mora, E. S.; Pal, U. Use of diffuse reflectancespectroscopy for optical characterization of un-supported nanostructures.ReV. Mex. Fis. 2007, 53 (5), 18.

(40) Nagaveni, K.; Hegde, M. S.; Ravishankar, N.; Subbanna, G. N.;Madras, G. Synthesis and structure of nanocrystalline TiO2 with lower bandgap showing high photocatalytic activity. Langmuir 2004, 20, 2900.

(41) Aarthi, T.; Madras, G. Photocatalytic degradation of rhodamine dyeswith nano-TiO2. Ind. Eng. Chem. Res. 2007, 46, 7.

(42) Sivalingam, G.; Nagaveni, K.; Hegde, M. S.; Madras, G. Photo-catalytic degradation of various dyes by combustion synthesized nanoanatase TiO2. Appl. Catal. B 2004, 48, 83.

(43) Linsebigler, A. L.; Lu, G.; Yates, J. T., Jr. Photocatalysis on TiO2

surfaces: Principles, mechanisms, and selected results. Chem. ReV. 1995,95, 735.

(44) Peng, T.; Zhao, D.; Dai, K.; Shi, W.; Hirao, K. Synthesis of titaniumdioxide nanoparticles with mesoporous anatase wall and high photocatalyticactivity. J. Phys. Chem. B 2005, 109, 4947.

(45) Yu, J; Zhang, L; Cheng, B; Su, Y. Hydrothermal Preparation andPhotocatalytic Activity of Hierarchically Sponge-like Macro-/MesoporousTitania. J. Phys. Chem. C 2007, 111, 10582.

(46) Rolison, D. R. Catalytic nanoarchitectures: The importance ofnothing and the unimportance of periodicity. Science 2003, 299, 1698.

(47) Wang, X; Yu, J. C; Ho, C; Hou, Y; Fu, X. Photocatalytic Activityof a Hierarchically Macro-/Mesoporous Titania. Langmuir 2005, 21, 2552.

ReceiVed for reView June 4, 2010ReVised manuscript receiVed August 20, 2010

Accepted September 11, 2010

IE101226B

Ind. Eng. Chem. Res., Vol. 49, No. 20, 2010 9643