Titanium dioxide mediated photocatalyzed degradation of a textile dye derivative, acid orange 8, in aqueous suspensions

Titanium dioxide mediated photocatalyzed degradation of a textile dye derivative, acid orange 8, in aqueous suspensions

~" ~ r~ DESALINATION ELSEVIER Desalination 155 (2003) 255-263 www.elsevier.com/locate/desal Titanium dioxide mediated photocatalyzed degradation...

679KB Sizes 1 Downloads 268 Views

~"

~

r~

DESALINATION ELSEVIER

Desalination 155 (2003) 255-263

www.elsevier.com/locate/desal

Titanium dioxide mediated photocatalyzed degradation of a textile dye derivative, acid orange 8, in aqueous suspensions Mohd Saquib, Mohd Muneer* Department of Chemistry, Aligarh Muslim University, Aligarh 202002, India Tel. +91 (571) 270-3515; Fax +91 (571) 270-2758; email: chtl2mm @ainu.hie.in

Received 8 January 2002; accepted 9 December 2002

Abstract

Titanium dioxide mediatedphotocatalyseddegradation of a textile dye derivative, acid orange 8 ( 1), was investigated in aqueous suspensions of titanium dioxide by mordtoting the depletion of total organic carbon (TOC) content as a function of irradiation time under a variety of conditions. The degradation kinetics were studied under different conditions such as pH, catalyst concentration, substrate concentration, different types of TiO2 and in the presence of electron aeceptors such as hydrogen peroxide (1-I20z) and potassium bromate (KBrO3) besides molecular oxygen. The degradation rates were found to be strongly influenced by all the above parameters. The photocatalyst Degussa P25 was found to be more efficient compared with other photocatalysts. The dye was found to be adsorbed on the surface of the photocatalyst at acidic pH. Keywords: Photocatalysis; Textile dye; Acid orange 8; Titanium dioxide

I. I n t r o d u c t i o n

The environmental problem associated with textile dyeing and finishing activities mainly originate from the extensive use of a wide variety of organic dyestuffs, the chemical characteristics of which vary greatly [1,2]. Azo dyes, characterized by the presence of the N=N linkage, comprise about half of all textile dyes used today [3]. One major source of these effluents is the *Corresponding author.

waste arising from industrial processes which utilize dyes to color paper, plastics and natural and artificial fibers [4]. A substantial amount of dyestuff is lost during the dyeing process in the textile industry [4], which poses a major problem for the industry as well as a threat to the environment [4-8]. Decolourization of dye effluents has therefore acquired increasing attention. Decolourization of dye effluent by bisulfite-catalysed borohydride reduction has also been reported [9]. During the past two decades, photocatalytic processes involving TiO2 semiconductor particles

0011-9164/03/$- See front matter © 2003 Elsevier Science B.V. All rights reserved PII: S 0 0 1 1 - 9 1 6 4 ( 0 3 )00303-5

256

M. Saquib, M. Muneer / Desalination 155 (2003) 255-263

under UV light illumination have been shown to be potentially advantageous and useful in the treatment of waste water pollutants. Earlier studies [ 10-13] have shown that a wide range of organic substrates such as alkanes, alkenes, aromatics, surfactants, and pesticides can be completely photomineralized in the presence of TiO 2 and oxygen. There are several studies related to the use of semiconductors in the photomineralization of photostable dyes [14-20]. Photoexcitation of semiconductors leads to the formation of an electron hole pair, which can eventually bring about a redox reaction from organic substrates dissolved in water. Alternatively, direct absorption of light by the dye, adsorbed on the semiconductor surfaces, can lead to charge injection from the excited state of the dye to the conduction band of the semiconductor. It has been shown earlier [ 14-20] that such processes can be used for removing colouring material from dye effluents in the presence of visible light. No major efforts have been made to study the degradation kinetics of an azo dye derivative such as acid orange 8. Thus, we have undertaken a detailed study on the photodegradation of this textile dye derivative, sensitized by TiO2 in aqueous solution.

/CH3 NaO3S ~

N--N

2. Experimental 2.1. Reagents and chemicals

Acid orange 8 (1) was obtained from Aldrich and used without any further purification. The water used in all the studies was double distilled. While the photocatalyst titanium dioxide, P25

(Degussa AG), was used in most of the experiments, other catalyst powders - - namely Hombikat UV100 (Sachtleben chemic GmbH), PC500 (Milenium inorganic chemicals) and TTP (Travancore titanium products, India) - - were used for comparative studies. P25 consists of 75% anatase and 25% rutile with a specific BETsurface area of 50 m2g-~ and a primary particle size of 20 nm [21]. Hombikat UV100 consists of 100% anatase with a specific BET-surface area >250 m2g-~ and a primary particle size of 5 nm [22]. The photocatalyst PC500 has a BET-surface area of 287 m2g-~ with 100% anatase and a primary particle size of 5-10 nm [23], whereas TTP has a BET-surface area of 9.82 mEg-1. The other chemicals used in this study such as NaOH, HNO3, H202 and KBrO 3, were obtained from Merck. 2.2. Procedure

Stock solutions of the dye containing the desired concentration were prepared in double distilled water. An immersion well photochemical reactor made of Pyrex glass equipped with a magnetic stirring bar, a water circulating jacket and an opening for supply of molecular oxygen was used. For irradiation experiments 250 mL of the stock solution were taken into the photoreactor and the required amount of photocatalyst was added; the solution was stirred and bubbled with molecular oxygen for at least 30 min in the dark to allow equilibration of the system so that the loss of compound due to adsorption could be taken into account. The zero time reading was obtained from a blank solution kept in the dark but otherwise treated similarly to the irradiated solution. The suspensions were continuously purged with molecular oxygen throughout each experiment. Irradiations were carried out using a 125 W medium-pressure mercury lamp. IR-radiation and short-wavelength UV-radiation were eliminated by a water jacket. Samples (10 mL) were collected before and at

M. Saquib, M. Muneer/ Desalination 155 (2003) 255-263 regular intervals during the irradiation. They were centrifuged by a Janetzki T-24 laboratory centrifuge before analysis.

2.3. Analysis The degradation was monitored by measuring the total organic carbon (TOC) content with a Shimadzu TOC 5000A analyzer by directly injecting the aqueous solution. For each experiment the degradation rate of the model pollutant was calculated from the initial slope obtained by linear regression from a plot of the natural logarithm of the TOC of the dye as a function of irradiation time, i.e., first-order degradation kinetics. It was calculated in terms of [mole Lmin-l].

3. Results

3.1. Photolysis of Ti02 suspensions containing acid orange 8 Fig. 1 shows the depletion in TOC on irradiation of an aqueous solution of acid orange 8 (0.25 mM) in the presence and absence of TiO2 by the Pyrex-filtered output of a 125 W medium-

257

pressure mercury lamp. It was found that 75% mineralization of the model compound took place after 90 min of illumination, whereas no observable loss of the dye occurred in the absence of the photocatalyst. The degradation curves can be fitted reasonably well by an exponential decay curve, suggesting first-order kinetics. The resulting first-order rate constant was used in all subsequent plots to calculate the degradation rate using the formula below: - d [TOC]/dt = kc" where k is the rate constant, c the concentration of the pollutant, and n is the order of reaction. Control experiments were carried out in all cases using unirradiated blank solutions. It has been found that the dye is adsorbed significantly in the dark at pH 3 and 5.3. The zero irradiation time readings were obtained from blank solutions kept in the dark, but otherwise treated similarly to the irradiated solutions.

3.2. Comparison of different photocatalysts The influence of four different photocatalysts (Degussa P25, Hombikat UV100, PC500 and TTP) on the degradation kinetics of acid orange 8 (1) is shown in Table 1. It was observed that

1,3. 1 -,,,,m--.~

I,-

rsoeoca~dya

Table 1 Comparison of degradation rate for the mineralization of acid orange 8 under different photocatalysts Type of photocatalyst

O,4

Degradation rate (mole L-1 min-j ×10-3)

02 0

i

o

i

I

:30

i

,

1 40

,

,

t !0

I

i

I

80

,

,

400

Irradiation time (rain)

Fig. 1. Influence of photocatalysts on the degradation of acid orange 8. Experimental conditions: dye concentration (0.25 mM), V = 250 mL, immersion well photoreactor, 125 W medium-pressure Hg lamp, photocatalyst: Degussa P25 (1 gL-~), cont. 02 purging and stirring, irradiation time = 1.5 h.

P25 0.00392 UV100 0.00265 PC500 0.00173 TTP 0.00043 Experimental conditions: dye concentration (0.25 mM), V= 250 mL, P25 (1 gL-l), Sachtleben Hombikat UV 100 (1 gL-t), PC 500 (1 gL-t), TTP (1 gL-~), immersion well photoreactor, 125W medium-pressure Hg lamp, cont. 02 purging, irradiation time = 1.5 h.

258

M. Saquib, M. Muneer/ Desalination 155 (2003) 255-263

4"00JOOS

~ ,'

i

i

I

'

'

'

I

'

'

'

I

'

'

'

I

'

'

'

I

'

'

'

' . 1.

.

'

.

I

'

'

'

'

I

'

'

'

'

I

'

'

'

'

I

'

I

'

'

I

'

'

'

,

'

1

"=n O~ iS

\

f,

-'~

pH-5.3



On04

a tm

-

i

o2 {} ,

0

,

,

l

2

,

,

,

I

4

,

,

,

I

6 pit

,

,

,

I

f

,

,

I

,

,

i

i

I

I

i

,

,

i

I

,

I

,

I

I

,

I

i

i

I

I

I

'

'

I

i

,

,

,

i

6

Fig. 2. Influence of pH on the degradation rate for the mineralizationof acid orange 8. Experimentalconditions as in Fig. 1. the degradation of dye proceeds much more rapidly in the presence of Degussa P25 compared with other photocatalysts. In all of the following experiments, Degussa P25 was used as the photocatalyst since this material exhibited the highest overall activity for the degradation of the model compound.

3.3. p H effect Using Degussa P25 as the photocatalyst, the photomineralization of acid orange 8 (1) in the aqueous suspensions of TiO 2 was studied in the pH range between 3 to 11. The degradation rate for the mineralization of acid orange 8 (1) as a function of reaction pH is shown in Fig. 2. The degradation rate for the TOC depletion is found to gradually increase from pH 3 to 9 followed by a sudden decrease as the pH rises to pH 11. A significant adsorption of the dye on the surface of the photocatalyst was observed at pH 3 and 5.3 as seen by the typical TOC V amount of TiO 2 at different pH (Fig. 3). The adsorption study ofazo dye 1 was investigated by stirring the aqueous solution in the dark for 24 h in a round-bottomed flask containing varying amounts of TiO2 such as 0,0.5, 1, 2 and 5 gL-] at pH 3, 5.3, 9 and 11. It is pertinent to mention here that the pH of the solution was adjusted before irradiation and it is not maintained throughout the reaction. A

O

1

2 3 4 Amou.t of TiO 2 (gL'l)

5

Fig. 3. Depletion in TOC Vs amount of the T i O 2 at different pH showing the adsorption of acid orange 8 in the dark. Experimental conditions: dye concentration (0.25 mM), V = 250 mL, Photocatalyst: Degussa P25 (0.5, 1, 2 and 5 gL-l), stirring in the dark for 24 h.

decrease in the pH of the reaction mixture was observed at the end of illumination.

3.4. Effect of substrate concentration It is important both from a mechanistic and an application point of view to study the dependence of photocatalytic reaction rates on the substrate concentration. Hence the effect of substrate concentration on the degradation of acid orange 8 (1) was studied at different concentrations: 0.12, 0.25, 0.35 and 0.5 mM. Fig. 4 shows the degradation rate for the mineralization of acid orange 8 as a function ofsubstrate concentration using Degussa P25 as the photocatalyst. It is interesting to note that the degradation rate increased with the increase in substrate concentration from 0.12 to 0.25 mM. A further increase in the substrate concentration from 0.25 to 0.5 mM led to a decrease in the degradation rate.

3.5. Effect of catalyst concentration The effect of photcatalyst concentration on the degradation kinetics of acid orange 8 (1) was investigated using different concentrations of Degussa P25 varying from 0.5 to 5 gL-L As

o •~

259

M. Saquib, M. Muneer / Desalination 155 (2003) 255-263

ao0s

'~D

0Din

|

00W

oom

'~ ooa ,

+ i I I , O1

I i I i i I , i , O2

, , ,

O3

i

~1

, ,

, , i O,5

,

k

*

*

06



0

+,,ill 1

Substrate Concent ration ( r a M )

Fig. 4. Influence of substrate concentration on the degradation rate for the mineralization of acid orange 8. Experimental conditions: dye concentrations (0.12, 0.25, 0.35, and 0.5 mM), V= 200 mL, Photocatalyst: Degussa P25 (1 gL-t), immersion well photoreactor, 125 W medium-pressure Hg lamp, cont. O5 purging and stirring, irradiation time = 1.5 h.

expected, the degradation rate of the acid orange 8 has been found to increase with the increase in catalyst concentration from 0.5 gLto 2 g L - I A further increase in the catalyst concentration from 2 to 5 gL- t leads to a decrease in the degradation rate (Fig. 5) due to the adsorption of the dye on the surface of the photocatalyst. This observation indicates that beyond this optimum concentration, other factors affect the degradation of model compound. At high TiO2 concentrations, particles aggregate, which reduces the interfacial area between the reaction solution and the catalyst; thus, they decrease the number of active sites on the surface. Light scattering by the particles and the increase in opacity may be other reasons for the decrease in the degradation rate. 3.6. Effect o f electron acceptors

Since hydroxyl radicals appear to play an important role in photocatalysis, electron acceptors such as hydrogen peroxide and potassium bromate were added to the solution in order to enhance the formation of hydroxyl radicals and also to inhibit the (e-/h +) pair recombination. The

2

3

Calalyst

concent~t~n

4

$

6

( g L "1)

Fig. 5. Influence of catalyst concentration on the degradation rate for the mineralization of acid orange 8 at different photocatalyst concentrations. Experimental conditions: dye concentration (0.25 mM), V= 250 mL. Photocatalyst: Degussa P25 (0.5, 1, 2 and 5 gL-l), immersion well photoreactor, 125 W medium-pressure Hg lamp, cont. O2 purging and stirring, irradiation time =1.5 h.

m

|

_tm J

}o

I~.~TH:O z

P2Y~'KBrO j

P25

Electron a c c e p t o r s

Fig. 6. Degradation rate for the mineralization of acid orange 8 in the presence of different electron acceptors. Experimental conditions: dye concentration (0.25 mM), V = 250 mL, P25 (1 gL-]), electron acceptors: KBrO3 (5 mM), HzO2 (10 mM), immersion well photoreactor, 125 W medium-pressure Hg lamp, cont. 02 purging and stirring, irradiation time = 1.5 h.

degradation rates for the mineralization of acid orange 8 in the presence of these electron acceptors are shown in Fig. 6. Both the additives such as hydrogen peroxide and potassium bromate had a beneficial effect on the degradation of the model compound.

260

M Saquib, M. Muneer / Desalination 155 (2003) 255-263

4. Discussion

The photocatalysed degradation of various organic systems using irradiated TiO2 is well documented [10]. The initial step in the TiO2 mediated photocatalysed degradation is proposed to involve the generation of the (e-/h ÷) pair leading to the formation of a hydroxyl radical and superoxide radical anion [Eqs. (1)-(3)]: TiO 2 + ho --~ e~b+ h~b

(1)

02 + e~b ~

(2)

H20 + h~b ~

0"2OH" + H ÷

(3)

It has been suggested that the hydroxyl radicals and superoxide radical anions are the primary oxidizing species in the photocatalytic oxidation processes. These oxidative reactions result in the bleaching of the dye, and the efficiency of the degradation depends upon the oxygen concentration, which determines the efficiency with which the conduction band electrons are scavenged and the (e-/h ÷) recombination is prevented. Alternatively, the electron in the conduction band can be picked up by the adsorbed dye molecules, leading to the formation of a dye radical anion; subsequent reaction of the radical anion can lead to degradation of the dye. In the present case both mechanisms can operate in an oxygen-saturated solution. The results on the photodegradation of the model compound using different kinds of TiO2 photocatalysts with different bulk and surface properties, i.e., BET-surface, impurities, lattice mismatches or density of hydroxyl groups on the catalyst's surface, are apparently responsible for the photocatalytic activity since they affect the adsorption behavior of a pollutant or intermediate molecule and the lifetime and recombination rate of electron-hole pairs. An earlier study [24] has shown that Degussa P25 owes its high photoreactivity to a slow

recombination of electron and holes whereas Sachtleben Hombikat UV 100 has a high photoreactivity due to its fast interfacial electron transfer rate. In these studies it was found that Degussa P25 showed better photocatalytic activity for the degradation of a model compound. Earlier studies have shown that Degussa P25 was found to show better activity for the photocatalytic degradation of a large number of organic compounds [25-28]. On the other hand, Lindner et al. [29] showed that Hombikat UV 100 was almost four times more effective than P25 when dichloroacetic acid was used as the model pollutant. The photocatalyst Hombikat UV 100 was found to be better for the degradation of benzidine and 1,2-diphenyl hydrazine as shown in a recent study [30]. These results indicate that the activity of the photocatalyst also depends on the type of the model pollutant. Another reason for the better efficiency of the Degussa P25 photocatalyst can be explained by the fact that Degussa P25 is a mixture of 25% rutile and 75% anatase. An important parameter in the photocatalytic reactions taking place on the particulate surfaces is the pH of the solution since it dictates the surface charge properties of the photocatalyst and size of aggregates it forms. For TiO2 the with Degussa P25 being the photocatalyst, the zero point of charge (pH~pc) is at pH 6.25. Hence, at more acidic pH values, the particle surface is positively charged, while at pH values above 6.25, it is negatively charged [31 ]. In this study it was shown that the degradation rate for a model compound under investigation is strongly influenced by the reaction pH, which increases with the increase in pH from 3 to 9 followed by a sudden decrease as the pH rises to 11. Due to the low pKa value of the sulphonic group, acid orange 8 was fully in the anionic form within the pH range studied. With increasing pH, the negative charges on TiO2 are expected to repel the dye, and a decrease in the efficiency of photo-degradation with increasing

M. Saquib, M. Muneer / Desalination 155 (2003) 255-263

pH is expected. However, it was observed that the rate increased with an increase in pH. Similar results were reported earlier for the acetate ion and acid blue 40 [32,33]. This effect may be attributed to more efficient generation of hydroxyl radicals by TiO2 with an increasing concentration of OH-. At the alkaline pH values, the hydroxyl radicals have to diffuse away and degrade the dye in the bulk solution. The dye was found to be significantly adsorbed in the dark at pH 3 and 5.3 as shown in Fig. 3. It is interesting to note that at pH 9 where the adsorption of the dye on the surface of the photocatalyst is negligible, the degradation rate for the mineralization of the dye is maximum, which suddenly decreases at pH 11. This effect can be attributed to the fact that the molecule may be undergoing structural changes in the double-bond character leading to the oxo form, which normally happens at higher pH values. The effect of substrate concentration on the degradation rate for the mineralization of 1 (Fig. 4) was studied, as it is important both from a mechanistic and the application point of view. As oxidation proceeds, less and less of the surface of the TiO2 particle is covered as the pollutant is decomposed. Evidently, at total decomposition, the rate of degradation is zero and a decreased photocatalytic rate is expected with increasing illumination time. It has been agreed, with minor variations, that the expression for the rate of photomineralization of organic substrates with irradiated TiO2 follows the Langmuir-Hinshelwood (L-H) law for the four possible situations, i.e., (a) the reaction takes place between two adsorbed substances, (b) the reaction occurs between a radical in solution and an adsorbed substrate molecule, (c) the reaction takes place between a radical linked to the surface and a substrate molecule in solution, and (d) the reaction occurs with both species being in solution. In all cases, the expression for the rate equation is similar to that derived from the L-H model, which has been useful in modeling the

261

process, although it is not possible to find out whether the process takes place on the surface, in the solution or at the interface. Our results on the effect of the initial concentration on the degradation rate of the model compound, shown in Fig. 4, indicate that the rate increases with the increase in the substrate concentration from 0.125 to 0.25 mM. A further increase in substrate concentration leads to a decrease in the degradation rate. This may be due to the fact that, as the initial concentrations of the dye increase, more and more dye molecules are adsorbed on the surface of the catalyst. Hence, the penetration of light to the surface of the catalyst decreases and the relative amount of OH" and 02- on the surface of the catalyst do not increase as the intensity of light and illumination time is constant. Conversely, their concentrations decrease with the increase in concentration of the dye as the light photons are largely absorbed and prevented from reaching the catalyst surface by the dye molecules. Consequently, the degradation efficiency of the dye decreases as the dye concentration increases. One practical problem in using TiO2 as a photcatalyst is the undesired electron/hole recombination, which, in the absence of proper electron acceptor or donor, is extremely efficient and thus represents the major energy-wasting step, which limits the achievable quantum yield. One strategy to inhibit the electron-hole pair recombination is to add other (irreversible) electron acceptors to the reaction, which could have several different effects, i.e., (1) to increase the number of trapped electrons and, consequently, avoid recombination, (2) to generate more radicals and other oxidizing species, (3) to increase the oxidation rate of intermediate compounds and (4) to avoid problems caused by low oxygen concentration. It is pertinent to mention here that in highly toxic wastewater where the degradation of organic pollutants is the major concern, the addition of inorganic ions to enhance the organic degradation rate may often be justified. In this connection, we

262

M. Saquib, M. Muneer / Desalination 155 (2003) 255-263

studied the effect of electron acceptors such as hydrogen peroxide and bromate on the photocatalytic degradation of the model compound under investigation. These acceptors are known to generate hydroxyl radicals according to the Eqs. (4)-(6):

H202 + ecB ~

O'H + OH-

BrO3 + 2H + + ecB

---}

BrO'2 + H20

BrO~ + 2H ÷ +6 eca -~ [BrO 2, HOBr] -~

(4) (5) (6)

BY + 3H20 As expected, both additives, hydrogen peroxide and potassium bromate, showed a beneficial effect on the photocatalytic degradation of model compound as shown in Fig. 6. The enhanced degradation rate in the presence of H202 can be explained by several reasons. Firstly, it increases the rate by removing the surface-trapped electrons, thereby lowering the electron-hole recombination rate and increasing the efficiency of hole utilization for reactions such as (OH- + h + --, O'H). Secondly, H202 may split photolytically to produce OH radicals directly, as cited in studies of homogeneous photooxidation using UV/(H202 + 02) [34]. Thirdly, the solution phase may at times be oxygen starved because of either oxygen consumption or slow oxygen mass transfer; peroxide addition thereby increases the rate towards what it would have been had an adequate oxygen supply been provided.

5. Conclusions T i O 2 c a n efficiently photocatalyse textile dye derivatives such as acid orange 8 using artificial radiation sources. The observations of these investigations clearly demonstrate the importance of choosing the optimum degradation parameters to obtain a high degradation rate, which is essential for any practical application of photo-

catalytic oxidation processes. The best degradation condition depends strongly on the kind of pollutant.

Acknowledgements Financial support by the Department of Science and Technology (DST), Government of India (New Delhi), and the Third World Academy of Sciences (Trieste, Italy) is gratefully acknowledged. The total organic carbon analyzer used for the analyses of the samples was a gift from the Alexander Von Humboldt Foundation (Germany).

References [ 1] A. Reifeand H.S. Freeman,Environmental Chemistry of Dyes and Pigments, Wiley, New York, 1996. [2] C. O'Neill, F.R Hawkes, D.L Hawkes, N.D Lourenco, H.M Pincheiro and W. Deice, J. Chem. Technol. Bioteehnol., 74 (1999) 1009. [3] J.R Easton, Colour in Dye,house Effluent, P. Cooper, ed., SDC, Bradford, 1995, p. 9. [4] H. Zollinger, in: Colour Chemistry, H.F. Eblel and C.D. Brenzinger,eds., 1st ed.,VCH,NewYork, 1987, Chap. 16. [5] C.E. Searle, Chemical Carcinogenesis, ACS Monograph, Amer. Chem. Sot., Washington, DC, 1976. [6] C.T. Helmes, C.C. Sigman, Z.A. Fund, M.K. Thompson, M.K. Voeltz, M. Makie, P.E. Klein and J.B. Lent, J. Environ. Sci. Health A, 19 (1984) 97. [7] M. Boeninger, Carcinogenicity and metabolism of azo dyes, especially those derived from benzidine, DHI-IS(NIOSH), PublicationNo. 80-119, July, 1980. [8] J.J. Roxon, A.J. Ryan and S.E. Wright, Food Cosmet. Toxicol., 5 (1967) 367. [9] M.M. Cook, J.A. Ulman and A. Iraelidis,Abstracts of Papers, 203 National Meeting of the Amer. Chem. Soc., San Francisco, 1992, and references therein. [10] D.M. Blake, Bibliography of work on the photocatalytic removal of hazardous compounds from water and air, National Renewal Energy Laboratory, USA, 1999. [11] J.M. Hermann, Catalysis Today, 53 (1999) 115-129.

M. Saquib, M. Muneer / Desalination 155 (2003) 255-263

[12] M.I. Litter, Applied Catalysis B: Environmental, 23 (1999) 89-114. [ 13] A. Fujishima, T.N. Rao and D.A. Tryk, J. Photochem. Photobiol. C: Photochem. Rev., 1 (2000) 1-21. [14] K. Vinodgopal, I. Bedja, S. Hotechandani and P. V. Kamat, Langmuir, 10 (1994) 1767-1771. [15] K. Vinodgopal and P. V. Kamat, J. Photochem. Photobiol. A: Chem., 83 (1994) 141-146. [16] A. Mills, A. Belghazi, R.H. Davies, D. Worsley and S. Morris, J. Photochem. Photobiol. A: Chem., 79 (1994) 131-139. [17] M. Vautier, C. Guillard and J.M. Hermann, J. Catal., 20 (2001) 46-59. [18] I. Arsalan, I.A. Balcioglu and D.W. Bahnemann, Dyes and Pigments, 47 (2000) 207-218. [19] S. Sakthivel, B. Neppolian, B. Arabindo, M. Palanichamy and Murugesan, Ind. J. Engineering Mat. Sci., 7 (2000) 87-93. [20] J. Zhao, T. Wu, K. WU, K. Oikawa, H. Hidaka and N. Serpone, Environ. Sci. Technol., 32 (1998) 23942400. [21] R.I. Bickley, T.G. Carreno, J.S. Lees, L. Palmisano and R.J.D. Tilley, J. Solid State Chem., 92 (1991) 178-190. [22] M. Lindner, D.W. Bahnemann, B. Hirthe and W.D. Griebler, J. Sol. Energy Eng., 119 (1997) 120-125. [23] S. Rauer, Untersunchung von kommerziell erhaltlichen Titandioxiden hinsichtlich ihrer photokatalytischen Aktivtat, Diplomarbeit, fachhochschule Hannover, Fachbereich Maschinenbau Vertiefung Umwelt-und Verfahrenstechnil, Hannover, 1998.

263

[24] S.T. Martin, H. Hermann, W. Choi and M.R. Hoffmann, J. Chem. Soc. Faraday Trans., 90 (1994) 33153322. [25] M. Muneer, J. Theurich and D. Bahnemann, J. Photochem. Photobioi. A: Chem., 143 (2001) 213219. [26] M. Muneer and D. Bahnemann, Water Sci. Technol., 144 (2001) 331-337. [27] M. Muneer, J. Theurich and D. Bahnemann, Res. Chem. Intermed., 25 (1999) 667-683. [28] M. Saquib and M. Muneer, Dyes and Pigments, 53 (2002) 237-249. [29] M. Lindner, D.W. Bahnemann, B. Hirthe and W.D. Griebler, Novel TiO2 powders as highly active photocatalysts, in: Solar Water Detoxification; Solar Engineering, W.B. Stine, T. Tanaka and D.E. Claridge, eds. ASME, New York, 1995, pp. 339-408. [30] M. Muneer, H.K. Singh and D. Bahnemann, Chemosphere, 49 (2002) 193-203. [31] H. Hiegendorff, Untersuchungen zur Bedeutung der Adsorption in der Photokatalyse,PhD Thesis, Department of Chemistry, University ofHannover, Hannover, Germany, 1996. [32] C. Koval, Proc. Symp. on Photoelectrochemistry, Electrochemical Society, Pennington, NJ, 1991. [33] M. Muneer, R. Philip and S. Das, Res. Chem. Intermed., 23 (1997) 233-246. [34] G.R. Peyton and W.H. Glaze, Environ. Sci. Technol., 22 (1988) 761-767.