heterogeneous photocatalysis in environmental remediation

Dev. Chem. Eng. Mineral Process., 8(5/6), pp.505-550, 2000.

Heterogeneous Photocatalysis in Environmental

Remedia tion

Dingwang Chen, M. Sivakumar and Ajay K. Ray* Department of Chemical and Environmental Engineering, The National

University of Singapore, 10 Kent Ridge Crescent, SINGAPORE I19260

Semiconductor-based photocatalytic processes have been studied for nearly 1.5 years

due to their many intriguing advantages in environmental remediation and other

areas. In this paper, underlying principles and mechanism of the photocatalytic

degradation process were discussed, followed by thermodynamics, kinetics, mass

transfer effects were discussed. The kinetics of photocatalytic degradation was

analysed with the aim of determining the favourable conditions to obtain high

quantum yield. Primary parameters that injluence the process such as catalyst

loading, dissolved oxygen, pH, temperature, light intensity, crystalline structure, etc.

have been discussed in detail. Different types of photocatalytic reactors, process

eflciencies and applications in environmental engineering as well as in other areas

have also been mentioned. The main barrier to the commercialization of the processes

is the low quantum yield. However, it is expected that large-scale applications could

be achieved with signijkant progress by improvement of process pelformance.

Introduction Water pollution is a major problem confronting us in the 21" century, particularly in

developing countries. People have now realized that water is no longer an endless

resource. As the global population grows, so does the need for clean water. With the

rapid development of science and technology, many industries, such as chemical,

petrochemical, pharmaceutical and mining require large quantities of water. Even

* Author for correspondence.

505

D. Chen, M. Sivakumar and A.K. Ray

ultrapure water is extensively needed in the pharmaceuticals, microelectronics and

semiconductor industries. Unfortunately, the discharged water from these industries is

often contaminated with toxic organic and inorganic compounds. In the recent two

decades, the pollutants, coming from a variety of industrial and agricultural activities,

are contaminating water and air to an unacceptable level all over the world.

Meanwhile, government regulations and public pressure no longer tolerate untreated

discharges. Hence, the need for innovative and effective water treatment processes for

both industry and human environment is extremely obvious.

There are mainly two purposes in water purification study. One is to reduce the

pollution levels in discharged streams to meet the government regulations. The other

is to purify water to ultrapure water, mainly for industries such as pharmaceuticals,

microelectronics and semiconductor. An ideal water treatment process should

completely mineralize all the toxic species present in the waste stream without

leaving behind any hazardous residues. It should also be cost-effective. At the current

state of development, none of the treatment technologies has reached this ideal

situation. There are a number of waste disposal methods currently in practice with

varying degrees of success. Figure 1 lists different wastewater treatment technologies

either currently available or in varied stage of development.

At present, the disposal of the bulk of the industrial wastes is based on the

processes developed on phase transfer principles [I] , even though none of them is

completely satisfactory. Air stripping is widely employed for the removal of volatile

organic contaminants in wastewater. But, the process just transfers the pollutants from

water phase to air phase rather than destroying them. Granular activated carbon

(GAC) adsorption is another commercialized process for water purification. However,

the spent carbon, on which pollutants are adsorbed, must then be disposed again. For

above reasons attention is being given to alternative destructive oxidation processes

for water treatment [2].

In contrast to the above non-destructive technologies, chlorination and ozonation

are two destructive technologies used in water industry. But, both use strong oxidants

that are seriously hazardous, therefore of undesirable nature. Furthermore, it has now

been proven that ozonation generates small levels of bromate ions, a recognized

Heterogeneous Photocatalysis in Environmental Remediation

.- I!!! m T

507

D. Chert, M. Sivakumar and A.K. Ray

cancer-causing agent [3]. Biodegradation of industrial wastes is still not common as

the degradation rates are usually very slow and at higher concentrations of organic

contaminants, it presents difficulties during the operation [4]. Besides, toxic organics

at times kill the active microorganisms that are intended to degrade them. Although

incineration of organic wastes has been widely used in developing countries, it

releases other toxic species into the air, such as dioxin and furan 151. In order to deal

with the ever-increasing degenerating environment and to address the problems

present in the commercial water treatment technologies, researchers and scientists

have been developing innovative treatment processes. A relatively new class of

technologies that shows promise for treating water contaminated with organic

compounds. These new technologies are known as advanced oxidation processes

(AOPs). H202/W, O m , Hz02/03/UV, TiOZ/W and V W processes are major

AOPs [ 61. In H202/W process, H202 is cleaved into hydroxyl radicals under the illumination

of W (mainly at 254 nm). These hydroxyl radicals attack and decompose the organic

compounds in aqueous solution by hydrogen abstraction, electrophilic addition and

electron transfer. The oxidation rate of contaminants in this process is usually limited

by the rate of formation of hydroxyl radicals, due to the rather small absorption cross

section of H2OZ at 254 nm, in particular, in the cases where organic substrates wili act

as inner filter [6] . Like other hydroxyl radical generating degradation processes.

03/w process can also oxidizes a great variety of organic compounds in wastewater.

In this process, aqueous systems saturated with ozone are irradiated with W light of

254 nm. Oxidative degradation rates are much higher than those observed in

experiments where either W light or ozone has been used separately. The most

serious and rather specific problem in the technical development of this process is the

low ozone solubility in water and consequent mass transfer limitations in

photoreactors [7]. 0 3 / H 2 0 2 / U V is quite similar to 0 3 I L T V process. This process is

enhanced by the photochemical generation of hydroxyl radicals by the addition of

hydrogen peroxide, due to the dominant production of hydroxyl radicals from H202.

The vacuum ultraviolet ( V W ) process uses W - C at 190 nm emitted by excimer

508


lamp as light source. V W photolysis

generation of hydroxyl radicals.

of H20 is the main and efficient means of the

... (1)

In general, the V W process is very simple and has the particular advantage that

no chemicals need to be added. However, only a few research groups, so far, are

active in this field due to the very limited availability of the excimer light sources [6] .

The ultimate goal of AOPs is complete mineralization of organic compounds to

carbon dioxide and mineral acids [8]. Among the APOs, heterogeneous photocatalysis

(Ti02/UV) has received considerable attention in the last 20 years as an alternative for

treating water polluted with toxic organic compounds and some metal ions. This

process is based on aqueous phase hydroxyl radical chemistry and couples lower

energy UV-A light with semiconductors acting as photocatalyst, and has emerged as a

viable alternative for solving environmental problems overcoming many of the

drawbacks that exist in traditional water treatment technologies. Photocatalysis

research is driven by legislation in industrialized countries that encourage water

purification (decontamination, detoxification, decolorization, deodorization) and

simultaneous contaminant destruction. Several books [8- 171 and review articles [6,

18-22] have been published in the last years reporting studies on photocatalytic

processes both in gas and liquid phases. Heterogeneous photocatalysis differs from

the other AOPs as it employs a reusable catalyst and there is no need to add any

additional oxidants. Its main advantages over the commercial water treatment

technologies include: (i) The oxidation is powerful and indiscriminate leading to the

mineralization of almost all-organic pollutants in wastewater. Even carbon

tetrachloride which was considered as hydroxyl radical resistant 1231 could also be

mineralized [24]. (ii) The process effluents (COz, H20, and mineral acid) are benign

to the environment, so it is called a “green technology”. (iii) The oxidant used in the

process is atmospheric oxygen, and therefore, in general, there is no need for

additional oxidizing agents. (iv) The catalysts are cheap, non-hazardous, stable and

reusable. (v) The light required to activate the catalyst is low energy UV-A, and

509


therefore, solar illumination is possible. (vi) The process is commercially comparable

to activated carbon adsorption for intermediate and large capacities [lo, 251.

Photoelectrochemistry received a great deal of attention after the discovery by

Fujishima and Honda [26] in 1972 that U V irradiated TiOz electrodes immersed in

water split water into hydrogen and oxygen gases. As investigation in this field

progressed, researchers found that the potential of photocatalytic processes, a subset

of photoelectrochemical phenomena, as an alternative to oxidize organic compounds

in water to carbon dioxide and mineral acids. Ollis and his co-workers [27-301 were

the first to implement photocatalysis as a method for water purification when they

conducted experiments for photomineralization of halogenerated hydrocarbon. Since

then, numerous studies have shown that a great variety of dissolved organic

compounds could be oxidized to CO, by photocatalysis [23, 31-36]. Apart from the

oxidation of organic compounds, reduction of some toxic metal ions [37-431 can also

be achieved by the heterogeneous photocatalysis. The overall process can be

described by the following two equations:

C,H,O, x + (m + 5 - 2 -2)o2 2 4 2

n-kz H20+kzH' +zX0Fk ... (2) Semiconduaor, hv ,mCO;? +- 2

> M o +nH' +LO2 ... Semiconduaor, hv M"' + n H 2 0 2 4 (3)

Basic Principles of Heterogeneous Photocatalysis The term photocatalysis consists of the combination of photochemistry and catalysis

and thus implies that light and catalyst are necessary to bring about or to accelerate a

chemical transformation [44, 451. The catalyst used are inevitably semiconductors,

which can act as catalyst due to their specific electronic structure characterized by a

filled valence band and an empty conduction band [46]. The energy gap, called

bandgap energy, between the conduction band and the valence band is relatively small

usually in the order of a few electron-volts. A photon with sufficient energy can

activate the catalys by promoting an electron from the filled valence band into the

conduction band. The occurrence of this charge separation is one of the first essential

510


steps of a photocatalytic reaction. This process is illustrated in Figure 2 and is

represented by Equation (4):

hV>E,,# TiO, , e , + h i ... (4)

0 2

k o x . 3

R' Minerals

A

OH'ads - o2 ___t Minerals RLJ kox.3

RHads

Figure 2. Schematic illustration of generation of electron hole pairs in a spherical semiconductor particle together with some of the consecutive redox reactions that take place.

511


As the conduction band is only partially filled, the electron remains free to move

through the semiconductor lattice. The resulting vacancy in the now partially filled

valence band is also free to move. This vacancy, or absence of an electron, is usually

referred to as hole, designated by h;. The photogenerated electron-hole pairs can

subsequently involve in two different reactions: (a) recombination of electrons and

holes and subsequent dissipation of the adsorbed energy with the liberation of heat,

and (b) reaction with electron donor or acceptor giving rise to oxidation and reduction

processes, respectively.

I A -1.0

0.0

w E 1.0

r w

C I

2.0

3.0

Figure 3. Valence and conductance band positions of semiconductors at pH = 0.

Figure 3 illustrates the valence and conductance band positions of various

semiconductors, and relevant redox couples. The position of conduction and valence

bands of a semiconductor determines the reaction pathway. For example,

photooxidation of water in presence of MoS2 (bandgap 1.8 eV) is possible

thermodynamically but is not possible in the presence of CdSe (bandgap 1.70 eV)

even though their bandgap energies are almost same. Instead, photo-reduction of

water would take place more easily in the presence of CdSe, due to the position of its

conduction band. Of course, more than one couple can be involved in a photocatalytic

512


process. Electrons and holes can reduce (and/or oxidize) species belonging to

different redox couples contemporaneously present in the reacting system. Hence, in

order for a semiconductor particle to be photochemically active as a catalyst for both

reactions (2) and (3), the redox potential of the photogenerated valence band hole

must be sufficiently positive to generate absorbed OH' radical (which is the powerful

oxidant for subsequent oxidation of organic pollutants), and the redox potential of the

photogenerated conductance band electron must be sufficiently negative to be able to

reduce the adsorbed oxygen to superoxide (or reduce the adsorbed metal ions, if there

is any). Figure 2 illustrates these reactions on the surface of a semiconductor, the

recombination process of the electron-hole pair, and some of the subsequent reactions.

The concentration of electron-hole pairs in a semiconductor particle is dependent on

the intensity of the incident light, and the material's electronic characteristics that

prevent them from recombining and releasing the absorbed energy.

Therefore, photocatalytic technology differs from other water treatment methods

in that it involves both oxidation and reduction chemistry. Reduction is mainly useful

for removing dissolved toxic metal ions from water whereas the oxidative property is

employed to mineralize dissolved toxic organic compounds. It should be emphasized

that oxidation and reduction must occur simultaneously to continue the process

activity. Organic compounds and water are two potential reductants, while oxygen

and metal ion are two potential oxidants in photocatalytic process. Oxidation and

reduction kinetics is interdependent, and either process can be rate-controlling 1471.

There are a variety of semiconducting materials which are commercially available

and investigated in literature as photocatalysts, such as TiOz, ZnO, ZnS, CdS, Fe203,

W03, etc. [48-52]. However, only a few of them were found out to be suitable for

efficiently catalyzing reactions in Equations (2) and (3) for a wide range of organic

and inorganic compounds. Table 1 lists the semiconductor photocatalysts commonly

used, their bandgap energy, and threshold wavelength as reported in literature. The

threshold wavelength is calculated using Equation (5):

Abg (nm> =124o/Ebg(eV) ... ( 5 )

However, the bandgap energies listed in Table 1 may be altered when the

semiconductor surface is in contact with an electrolyte solution [14, 531. Lower

513


bandgap energy semiconductors are preferred when catalyst activation with solar

energy is desired. However, low bandgap semiconductors, especially non-oxides such

as CdS and CdSe, usually suffer serious stability problem [54]. A desired

photocatalyst must posses: (a) high activity; (b) ability to utilize visible and/or near-

W light; (c) photostable (durable) and reusable; (d) chemically and biologically

inert; and (e) cheap. Of all the different semiconductor photocatalysts tested

successfully in laboratory studies. Ti02 appears the most active and is extensively

used in both laboratory studies and pilot plants [48, 55, 561. Ti02 is cheap, insoluble

under most conditions, photostable, and non-toxic. Moreover, sunlight (about 3% of

the solar spectrum contains W - A ) can be used as a possible light source for Ti02

(absorbs light in the UV-A range, h 5388 nm) activation [57-591.

Table 1. Bandgap energies (at pH = 0) and corresponding threshold wavelengths of various semiconductors [ I 61.

Semiconductor Bandgap (eV) Wavelength (nm) Ti02 3.0-3.2 4 13-388 ZnO 3.2 388 ZnS 3.6 335 CdS 2.4 516

FQ.03 2.3 539 wo3 2.8 443

Titanium dioxide exists primarily in two crystalline forms, anatase and rutile. In

many studies, anatase was found more effective than mi le for the photocatalytic

oxidation of acid/acetate mixture [60], phenol [61, 621, cyclohexane and 2-propanol

[63]. The anatase form has a reasonably well-defined nature, typically 70 to 30

anatase to rutile mixture, non-porous with BET surface area of 55+15 m2/g and

average particle size of 30 nm. It shows substantially higher photocatalytic activity

than most other readily available samples of TiO2.

Mechanism of the Photocatalytic Reaction Although Equations (2) and (3) represent the overall reactions, individual steps that

drive above reactions are still not well understood. The photocatalytic process can be

referred to as “four-phase” system [64]. The fourth phase being the electronic phase

(W light source), in addition to the three phases - liquid (water), solid (catalyst) and

514


gas (the oxidant, oxygen). The basic phenomena that take place in a semiconductor

particle in aqueous solution irradiated with ultraviolet light of appropriate wavelength

are schematically shown in Figure 2. When a semiconductor is illuminated by light of

suitable energy, namely higher than its bandgap energy Ebg, electrons are promoted

from the valence band to the conduction band, and consequently, the positive holes

are created in the valence band, which is represented by Equation (4). The electron-

hole generation is one of the first essential steps of a photocatalytic reaction. While in

contact with a liquid solution, a charge transfer occurs until an electrostatic

equilibrium is achieved. The excess of charge in the semiconductor is not confined on

the surface as occurs in a metal, but is distributed in a region called the 'space-charge-

region' expanding from the surface to the bulk of the semiconductor. This space-

charge-region plays a significant role in photocatalytic process. In the presence of the

surface charge region, the photo-excited electron-hole pair can be separated, and then

migrate ( kJ# , k, ) towards the particle surface (Equations 5 and 6) as a result of a

potential gradient that exists between the bulk solid and its external surface. This

potential gradient is caused by depletion of conduction band electrons in the space-

charge-region [ 19, 65, 661. Electro-neutrality of the surface requires equal arrival

rates of the electrons and holes at steady state [67]. The fate and dynamics of the

photogenerated electron-hole pairs are of considerably importance in the subsequent

degradation processes. Generally, the mobile electors and holes have two basic

destinies. They may recombine (k,) and dissipate the inputted energy as heat

(Equation 7) during their transport to the particle surface. They can also react with the

electron acceptor and electron donor on the particle surface after migration to the

catalyst surface, and subsequently, initiate the reduction and oxidation processes,

respectively (Equations 8 and 9).

- h+

e, + hv+b krec >heat ... (8)

51.5


A + e , kred 9 A'- ... (9)

Reaction (7) is one of the main causes for the low quantum efficiencies of

photocatalytic processes, and the recombination reaction of electrons and holes occurs

mainly in the bulk of the catalytic particle [68]. This recombination possibility can be

reduced if these mobile species are separated, and subsequently, trapped by surface

adsorbates or other sites. After a photogenerated hole reaches the surface of the

semiconductor, it can react with an adsorbed substrate by interfacial electron transfer,

assuming that the adsorbate possesses a redox potential appropriate for a

thermodynamically allowed reaction. Thus an adsorbed electron donor can be

oxidized by transfemng an electron to a hole on the surface, and an adsorbed acceptor

can be reduced by accepting an electron from the surface. Hole trapping generates a

cation radical Do+, and electron trapping generates an anion radical, A*-. These radical

ions can participate in several reaction pathways: (i) they may react chemically with

themselves or other adsorbates; (ii) they may recombine by back electron transfer to

form an excited state of one of the reactants, or to waste the excitation energy by the

non-radiative pathway; (iii) they may diffuse from the semiconductor surface, and

participate in chemical reaction in the bulk solution. If the rate of formation of D" is

kinetically competitive with the rate of back electron transfer, photoinduced oxidation

will occur for any molecule with an oxidation potential less positive than the

semiconductor valence band edge, since under these conditions interfacial electron

transfer at the illuminated interface is thermodynamically possible. For similar

considerations, the photoinduced reduction can occur to any molecule possessing a

reduction potential less negative than the conduction band edge [ 191.

It is well known that the surface of TiOz is readily hydroxylated when it is in

contact with water. Both dissociated and molecular water is bonded to the surface of

TiOz. Surface coverage of 7-10 OR/nm2 at room temperature were reported in

literature [69, 701. The reactions of the trapped holes react with absorbed solvent

molecules (H20 and OH) have been experimentally observed [32,71-731:

516


>OH>s + H+ ... (11)

,OH23 ... (12)

kox, 1

kox, 1

G + H20ads

h i + OHSs

The trapped holes can also react with absorbed organic substrates RX:

h: + J?X, *RXzs ... (13) Reactions (10) and (11) is more important in oxidative processes due to the high

concentration of the absorbed H2O and O H molecules on catalyst particle surface.

Experiments conducted in water-free, aerated organic solvents have shown that only

partial oxidation can be achieved "731. However, in aqueous solutions complete

mineralization of numerous organic compounds has been observed [27-29, 74-76].

Apparently, direct reaction between the organics and the valence band holes

(Equation 12) is not significant [77]. The result also implies that the presence of water

or hydroxyl group is necessary for complete destruction of organic compounds.

Photogenerated electron must also react to avoid a continuous charge build-up in

catalyst particles. The accumulation of photogenerated electrons on Ti02 particles

was experimentally observed during the photoassisted oxidation of 1.6M aqueous

methanol [77]. At steady state the rate of hole consumption must be equal to the rate

of electron consumption. Therefore, electron scavenger (or acceptor) must be present

in the photocatalytic process. Oxygen is the commonly used electron acceptor as it is

available at little or no cost. It reacts with photo-generated electron at the surface of

semiconductor through following equations [60, 61,73,78]:

e,; + 0 2 , udc * 0 : i ... (14)

H' + O;;& - HO; ... (15)

2HO; - H 2 0 , + 0, ... (16)

HO; + O;;ab - 0, + HO; ... (17)

H O ; + H ' - H , O , ... (18)

H , O , - H, + 0,

H , O , + H O ~ ~ O H * + H , O + O , ... (20)

... (19)

517

D. Chen, M. Sivakumar and A. K. Ray

Formation of these species has been experimentally verified by electron spin

resonance measurements [71]. Auguliaro et al. [79] also confirmed the formation of

H202 in the degradation of phenol in Ti02 suspensions by aerating the solution with

O2 and He respectively. Oxygen-involved species (e.g., HO;, HO;, 02'-, O2 and

H202) are present either at the interface or in the solution, and can participate in the

complex degradation scheme, which leads to the final mineralization of the organic

species. It should be noted that it takes three electrons to make one hydroxyl radical

through this pathway, but it takes only one hole to produce a hydroxyl radical in the

other half-cell reaction, Equations (10) and (1 1). Consequently, most hydroxyl

radicals are generated through the hole reaction [61, 801.

The hydrogen peroxide formed from reactions (13) to (15) may further decompose

to form hydroxyl radicals [61, 79, 811:

H20,*20H' ... (21)

H,O, + 0;- -OH + OH- + 0, ... (22)

H,O,+e,--+OH'+OH- ... (23)

It has been shown that the addition of hydrogen peroxide considerably improves

the photocatalytic process [36, 79, 82-86], most probably through reactions (20) to

(22). It has been proposed that the highly reactive hydroxyl radicals formed in above

equations are the primary oxidizing species in photocatalytic processes by attacking

the organic species and the rate-determining reaction step may be the formation of

hydroxyl radicals [32, 65, 731. This is supported by the following facts: (a) Complete

mineralization of organic compounds cannot be achieved in water-free organic

solvents [73], however it is possible in aqueous solutions; (b) Steady-state OH'

concentrations in W-irradiated titanium dioxide aqueous solutions could be as high

as lo-' M [87] much higher than those in the processes of ozonation, direct photolysis

of H202, and radiolysis (i.e., c 10l2 M); (c) During the photocatalytic degradation of

aromatic compounds, the detected intermediates typically have hydroxylated

structures [32, 61, 621. These intermediates are consistent with those found when

similar aromatics are reacted with a known source of hydroxyl radicals [88. 891.

518


Hydroxyl radical is a very reactive species with an unpaired electron. It reacts

rapidly and non-selectively in the oxidation of organic compounds. Due to its very

high oxidation potential (see Table 2), it is capable of oxidizing almost all organic

pollutants in wastewater by hydrogen abstraction. The reactions may occur either in

the bulk solution or on the catalyst surface. Based on hydroxyl radical attack ( kox, )

of organic compound in photocatalytic degradation, Turchi and Ollis [32] proposed

four different reaction pathways:

Reaction occurs while both species are adsorbed.

OHh + ' , , a h -';.ads

A non-bound radical reacts with an adsorbed organic species.

OH + 4.d -';,ah

An adsorbed radical reacts with a free organic species arriving at the

catalyst surface.

OH; + R, d R;

Reaction occurs between two free species in the bulk solution.

OH' + R, - R;

. . . (24)

... (25)

... (26)

... (27)

Based on the above different reaction pathways, the four kinetic models developed

had similar degradation rate forms and were consistent with reported initial data and

temporal degradation data. However, pulse radiolysis and time-resolved diffusion

reflectance measurements showed that surface-bound hydroxyl radicals are more

stable compared to those in the bulk solution [90]. Similar results were obtained by

Minero et al. [91] who found that the degradation process occurs at the surface or

within a few monolayers around the photocatalytic particles. These results imply that

the main reaction of the photocatalytic degradation process takes place on the surface

of the catalyst.

The organic radicals formed in above reactions can be oxidized ( kox, ) into

peroxy radical by addition of molecular oxygen. These intermediates initiate thermal

(chain) reactions leading to COz, H 2 0 and mineral acid:

5-19


RLfds + 0 2 RO; ++ C02 -i- H20 i- Mineral Acid ... (28)

Table 2. Oxidation Potentials of Some Oxidants 161. Species Oxidation Potential (V) fluorine 3.03

hydroxyl radical 2.80 atomic oxygen 2.42

ozone 2.07 hydrogen peroxide 1.78 perhydroxyl radical 1.70

permanganate 1.68 hypobromous acid 1.59

clorine dioxide 1.57 hypochlorous acid 1.49 hypoiodous acid 1.45

chlorine 1.36 bromine 1.09 iodine 0.54

Table 3. Reaction rate constants of hydroxyl radical with classes of organic contaminants.

Compounds Reaction rate constant (VmoVs) Chlorinated alkenes 1 09- 10'

Phenols 1 09- 10" N-containing organics lo9- lo1'

Aromatics 1 oE- 1 O'O Ketones 1 09- 1 o10

Alcohols 1 08-109 Alkanes lo6- lo9

Table 3 presents a summary of the classes of organic contaminants typically

treatable by advanced oxidation process (AOP) technology, as well as their associated

reaction rate constants relative to oxidation by hydroxyl radicals. The data indicate

that chlorinated alkenes, phenols and nitrogen-containing organics exhibit the highest

rate constants principally due to the presence of double bounds within the molecular

chain and other characteristics that make theses compounds susceptible to oxidation.

During photocatalytic oxidation of organic compounds, oxygen is reduced to

superoxide and/or perhydroxyl radicals. These radicals eventually form hydroxyl

radicals that enter into the oxidation cycle [3, 19, 201. Besides oxygen, any dissolved

metal ions with a reduction potential more positive than the conduction band of the

photocatalyst can, in principle, consume electrons and complete the redox cycle. This

520


usually deposits a reduced form of the metal on the catalyst surface, which may or

may not be easily removed [92, 931. Figure 4 illustrates the positions of valence and

conduction bands of anatase Ti02 photocatalyst in contact with an aqueous electrolyte

solution at different pH and compares them with the reduction potentials of

environmentally fretful metal ions. It is worth noticing that the positions of both

conduction and valence bands are pH dependent. The increase of pH in electrolyte

solution makes the positions of the valence and conduction bands to shift to more

cathodic potentials by 59mV per pH unit [47]. However, reduction potentials of all

metal ions (except for Cr(V1)) illustrated in Figure 4 are independent of pH, thereby

resulting in photoreduction of these metal ions more favourable with increasing pH.

On the basis of Figure 4, one can note that Cd2+, Fez+ and Cr3+ cannot be

photocatalytically reduced because their reduction potentials are more negative than

that of electrons. However, photocatalytic reduction is thermodynamically feasible for

Au3+, Cr", Hg2+ (including HgC12, HgCb'.), Ag', Hg?, Fe3+, Cu' and Cu2+. Among

the above metal ions, Fe3+ and Cr6+ can only be reduced to Fe2+ and Cr3+, rather than

iron and chromium metals respectively, because Fez+ and Cr3' cannot be reduced

further as mentioned above. It is also unlikely for Pb2+ and Ni2' to be reduced under

most conditions due to the extremely low driving force. One must also note that the

reduction potential of a redox couple is concentration dependent. According to Figure

4, one may also expect that oxygen can be reduced preferentially with respect to most

metal ions if it is present in solution. That is why the reaction system in most studies

was purged with nitrogen or sealed to eliminate oxygen reduction, thereby increasing

the photoreduction efficiency of metal ions. However, the photocatalytic reduction

rate or to what extent they can be removed for those metal ions that can be reduced

thermodynamically at given illumination time is determined by the reduction kinetics.

Sometimes, no measurable reduction is observed due to the very low reduction rate

even though the reduction reaction is thermodynamically feasible.

Photocatalytic Kinetics Unlike other AOPs that are based on the use of additional chemical reagents, the

kinetic behavior for chemical transformations in photocatalysis usually follows a

saturation behavior [94]. Therefore, typical applications of photocatalysis are at low

521


level of pollutant concentration. The kinetic modeling of the primary steps is required

for any practical application of the process and some kinetic models describing

photocatalytic oxidation on illuminated semiconductors have been proposed.

3.2

2.8 - - 1 -

I

Qz 2.4

1.2

0.8

Potential (V)

0.4

0

-0.4

-0.8 21

I I I I I I

0 1 2 3 4 5 6 7

UH

Figure 4. Positions of valence and conduction bands of Ti02 (anatase) and the reduction potentials of metallic ions of interest at different p H .

522


The degradation kmetic rate expressions in literature have focused on the initial

disappearance rate of organic compounds [e.g. 78, 961 or the initial formation rate of

C 0 2 [e.g. 33, 54, 971. The initial photocatalytic reaction rate usually follows

Langmuir-Hinshelwood equation with respect to initial organic concentration [e.g.

27-33, 62,94-97]:

kKCo ro =

1 + KC, .. (29)

The above relationship describes a zero-order rate at high concentration while a first

order reaction rate at low organic concentration. Figure 5 illustrates the typical

laboratory experimental set-up for photocatalytic kinetic studies. Usually,

experiments are carried out in a semi-batch mode, which mainly consists of

photoreactor, UV light source, pump, well-mixed reservoir and numerous measuring

meters. Adjustments can be easily performed in the reservoir for sampling and

suspension chemistry monitoring.

a TiOz suspension

W lamp

lol Magnetic stirrer

Figure 5. Typical experimental setup for photocatalytic kinetic studies in laboratory.

523

D. Chen, M. Sivakumr andA.K. Ray

Mass Transfer in Photocatalytic Processes In heterogeneous catalysis, both internal and external mass transfer of the reactants

and products affect the overall rate of degradation. Relationship among the observed

degradation rate, kobs, the external and internal mass transfer rates, km,e,, and km,int, and

the intrinsic kinetic rate, kn, are given by the following expression:

... (30) 1 1 1 - 1 ---+-+-

kobs krxn km,ext km,int

Therefore, the apparent rate of photocatalytic reactions can be either be reaction rate

controlled or mass transfer controlled.

Photocatalytic processes are generally operated in two modes where catalyst is

either suspended or immobilized. Although Degussa P25 TiOz is non-porous and its

elementary particle size is only about 30 nm, actual catalyst particle size of up to 5

pm in diameter is reported [98] due to the aggregation of the elementary particles

when suspended in aqueous solution. Chen and Ray [99] investigated the mass

transfer effect in P25 TiOz suspensions using phenol, 4-chlorophenol and 4-

nitrophenol as model compounds. The results indicated that ratios of surface

concentration to the bulk concentration were close to unity, and the minimum

effectiveness factors of actual catalyst particles were above 0.9. Therefore, in TiOz

suspensions both internal and external mass transfer resistance may be neglected and

chemical reaction is the rate determining (controlling) step. However, effect of mass

can not be ignored when catalyst is immobilized. Observed photocatalytic degradation

rate increases with increasing flowrate of reactants as reported [95, 1001 in both

straight and coiled tube photoreactor, where catalyst was coated on the inside surface

of the tube. Scalfani et al. [loll found that mass transfer was the main rate-limiting

factor in photodegradation of phenol in a continuous annular photoreactor in which

Ti02 beads were used. Ray and Beenackers [lo21 investigated the mass transfer in a

swirl-flow monolithic-type photoreactor by measuring the dissolution rate of benzoic

acid into water flowing at different Reynolds number. The mass transfer coefficient

was correlated by:

k, = 6.1~10-' ... (31)

524


When the thickness of catalyst layer (film) is increased, internal mass transfer may

also play a significant role in photocatalytic reactions. This influence was determined

by Chen et al. [I031 experimentally by conducting dynamic physical adsorption of

benzoic acid on P25 Ti02 catalyst film at different catalyst layer thickness. They

reported as effective diffusivity value of ~ . O X I O ' ~ m2/s by fitting the experimental

results with a realistic dynamic physical adsorption model.

Kinetic Equations Many researchers have demonstrated that variety of pollutants can be effectively

degraded using Ti02 as catalyst. However, the kinetics and mechanism of

photomineralization processes are still comparatively ambiguous. Several kinetic

models [e.g., 27, 32, 33, 54, 72, 75, 79, 80, 95, 104-1061 for photocatalytic oxidation

of organic compounds have been published in literature. Among these reaction

models, Langmuir-Hinshelwood type model was the most popular in describing the

photocatalytic oxidation, which was further divided into competitive and non-

competitive type adsorption models [32, 1071. The competitive model described that

both oxygen and organic compound adsorbed on the same active site of catalyst

competitively, while the non-competitive model specify that the organic compound

and oxygen adsorbed separately on the different active sites of the catalyst. Turchi

and Ollis [32] proposed that the non-competitive model was favored for the

photocatalytic oxidation using Ti02 as catalyst. Matthews [33, 75, 95, 1081 studied

the photocatalytic oxidation of organic substrates over TiOz, Okamoto [go]

investigated the photodecomposition of phenol, Pruden and Ollis [27] studied the

photodegradation of trichloroethylene in water, Chen and Chou [ 1051 investigated the

photodecolorization of methyl orange, and Mills and Morris [54] studied

photooxidation of 4-CP. All of these researchers have pointed out that the Langmuir

adsorption model could be applied to their systems as summarized in Table 4.

To date, very few photocatalytic kinetics of Ti02 on immobilized system have

been published. Sabate et al. [ 1061 investigated the photocatalytic degradation of 3-

chlorosalicylic acid over Ti02 film supported on glass. In their experiments, they

varied the concentrations of organic compound and dissolved oxygen, but no

525

Tabl

e 4. S

urve

y of

typi

cal k

inet

ic m

odel

s pro

pose

d in

lite

ratu

re

Pollu

tant

Ca

taly

st O

pera

ting m

ode

Rat

e uni

t K

inet

ic m

odel

Re

f. O

xalic

acid

Fi

xed

Batc

h (r0

) m

olh

kKi IS

1 55

Benz

oic a

cid,

etc*

CP

4-C

P

Tric

hlor

oeth

ylen

e

3-C

hlor

osal

icyl

ic ac

id

Chl

orof

orm

, etc

Phen

ol

Met

hyl o

rang

e

4-Ch

loro

phen

ol'

4-ni

troph

enol

Phen

ol

Slur

ry

Fixe

d

Slur

ry

Slur

ry

Fixe

d

Slur

ry

Slur

ry

Slur

ry

Slur

ry

Slur

ry

Slur

ry

Sem

i-bat

ch

Sem

i-bat

ch

Batc

h

Batc

h

Sem

i-bat

ch

batch

Batc

h

Batc

h

Batc

h

Batc

h

mgl

min

.L

moU

min

.L

mm

ol/h

.L

pmoU

min

.g.c

at

m0V

min

.L

mm

oUm

in

m0V

min

.L

mg/

min

.L

moV

m2.

s

rn0V

s.L

33, I

5

95

104

21

106

32

80

105

I40

56

19

P 9 P

P

a


influence of other parameters, such as catalyst loading, light intensity and reaction

temperature was reported. Ray and Beenackers [lo21 performed a detailed study for

degradation of a textile dye on immobilized system. They designed a novel kinetic

reactor using which they determined lunetic rate constants as a function of

wavelength of light intensity and angle of incidence, catalyst layer thickness, and the

effect of absorption of light by liquid film on the overall photocatalytic degradation.

Unfortunately, there are two main disadvantages in the reported kinetic equations.

Firstly, most of these equations either describe the relation between inirial

degradation rate and initial concentration of organic species, or used the kinetic

parameters obtained from the initial rate to predict the entire process. Therefore, good

agreement can be obtained only for those systems in which intermediates are not

formed during photocatalytic degradation such as PCE [31]. However, for most

photocatalytic processes, such as the photodegradation of halogenated aromatics,

above kinetic equations cannot predict the entire process due to the formation and

influence of the intermediates. Furthermore, initial rate data are tedious to obtain, and

prone to large errors, and thereby reduces reliability on the results. Secondly, in

heterogeneous catalysis, it is customary to report the reaction rates in units of per

gram of catalyst. However, heterogeneous photocatalysis is quite different from the

ordinary heterogeneous catalytic process due to the presence of W light. Catalyst at

different positions in a reactor has different contributions to the reaction rate due to

the light distribution within the reactor. In Ti02 aqueous suspensions, the catalyst that

is far from the light source has less or even no contribution to the reaction rate. When

TiOz is immobilized onto a substrate, not all the catalyst particles take part in the

reaction. Obviously, in heterogeneous photocatalytic reaction, it is not reasonable to

define the unit of rate as the same as those commensurate with general homogeneous

or heterogeneous phase kinetics. However, the definition of reaction rate reported in

the literature is either in “moVs” or “moWs” [33, 74, 78, 79, 96, 105, 109, 1101.

Since both are dependent on the illuminated catalyst surface area through the reactor

window and volume of reaction solution, the reported results are in general not

meaningful, and moreover, render the comparison between investigations conducted

in different research groups impossible. In order to avoid above problems, Chen and

527


Ray [56, 991 developed new photocatalytic kinetic expressions. Their model

correlated as many kinetic parameters as possible. Reaction rate was defined as the

mole reduction of pollutant per irradiated reactor window area. So, it was independent

of either the illuminated area or the volume of the reaction solution.

Kinetic Parameters Initial Concentration of organic compounds Influence of initial concentration of organic compounds on the photocatalytic

degradation rate has been extensively studied and the so-called "saturation behaviour"

was observed by many researchers [49,54,94,96, 105, 109, 11 11. Langmuir-type rate

expression (Equation 28) has been used successfully to describe photocatalytic

degradation. However, the mechanism of this effect is still not clear. We believe that

three factors might be responsible for the saturation behaviour during the

photocatalytic reaction. (a) One of the main steps in the degradation process take

place on the surface of the catalyst, and therefore, adsorption of organic compounds

onto the catalyst surface affects the reaction, and usually high adsorption capacity

favours the reaction. For most organic compounds adsorption capacities on TiOz

catalyst are well described by Langmuir-type equation. This has also been confirmed

in our laboratory study [59]. It means that at high initial concentration all accessible

catalytic sites are occupied. A further increase in pollutant concentration does not

increase concentration of pollutant at the catalyst surface. (b) In photocatalytic

processes, generation and migration of the photogenerated electron-hole pair, and the

reaction between photogenerated hole (or hydroxyl radical) and organic compound

are two processes in series. Hence, each step may become rate determining for the

entire process. At low organic concentration the latter may dominate the process and,

therefore, the degradation rate increases linearly with the concentration. On the

contrary, at high organic concentration the former will become the governing step and

the degradation rate increases slowly with concentration, and even a constant

degradation rate may be observed at higher concentration under a given illuminating

light intensity. Gerischer and Heller [81] and Wang et al. [77] provided substantial

evidence indicating that the interfacial electron transfer process involving the

reduction of oxygen was the rate determining step in the TiOz sensitized

528


photodegradation of organics. (c) Intermediates formed during photocatalytic process

also affect the rate constant of their parent compounds. For example, consider two

experimental runs with different initial concentrations CI and C2 (Cl > C2). When C, decreases to C2. some intermediates will be formed and subsequently will adsorb

competitiveIy on the solid catalyst surface. If the two experimental runs are now

compared starting with C1, the adsorption of intermediates will reduce the effective

concentration of the parent compound on the catalyst surface for first case over the

second. Hence, the observed rate for the first case will be lower.

Catalyst Dosage

In the photocatalytic oxidation of organic compounds in aqueous suspensions of Ti02,

the mass concentration of the latter is an important process parameter affecting the

observed reaction rates [27, 28, 105, 112, 1131. Most commonly, 0.1-0.2 wt% TiOz

slurry are used [36, 50, 75, 1141. In the studies by Ollis and co-workers [36], a 0.1

wt% Ti02 slurry was used to photocatalytically oxidize a variety of halogenated

organic compounds in a batch reactor [27-301. Okamoto et al. [61, 801 studied among

other variables the effect of Ti02 concentration in the photooxidation of phenol in

aqueous suspensions and found that at high light intensities the reaction rate increased

with the Ti02 concentration in the solution. Similar studies showed that an optimal

concentration range for Ti02 (1-3 g/L) exist for maximum removal of phenol 62.

1151. This optimum concentration range depends on the reactor geometry, intensity of

radiation source, and the properties of Ti02 such as particle size, phase composition

and impurities. Increasing catalyst concentration beyond the optimum range may

result in reduced photon flux caused by a shielding effect of the particles on the light,

although the catalyst surface area per unit volume of solution increases [lo, 62, 1151.

This shielding effect has been studied in a stirred photoreactor [ 1161. Table 5 lists the

reported research results on catalyst dosage used. Obviously, the optimal catalyst

dosage reported was in a wide range from 0.15 to 8 g L for different photocatalyzed

systems and photoreactors. Even for the same catalyst (Degussa P25), a large

difference in optimal catalyst dosage (from 0.15 to 2.5 g/L) was reported.

Chen and Ray [99] proposed the following equation that describes the influence of

catalyst dosage on the photocatalytic degradation rate:

529


ri = K[1-exp(+BCCatH)] ... (32)

where E is the light adsorption coefficient of the reaction system; p is the exponential

term of light intensity influence and its value is between 0.5 and 1 .O depending on the

light intensity; H is solution thickness in light transmission direction. Equation (3 1)

can successfully explain the dependence of degradation rate on the Ti02 dosage at

low and high light intensities reported in literature [80]. Degradation rate is usually

proportional to between p5 and I].' at high and low intensities, respectively [ 18, 36,

94, 1021. This indicates that the p value is 0.5 at high intensity and 1.0 at low

intensity. Thus, according to Equation (31) the optimal catalyst dosage at high light

intensity is higher than that at low light intensity. Experiments conducted under

different solution thickness indicated that effective optical penetration length was

only a few centimeters in 0.2% Ti02 suspensions [99].

Table 5. List of optimal Ti02 catalyst dosage reported in literature.

Ref. Optimal dosage (g/L) Pollutant(s) Catalyst

2- and 3-chlorophenol Degussa P25 2.5 94

4-chlorophenol Degussa P25 2.0 50 Nitrophenols BDH 1 .o 49

Phenol Merck 2.5 115

Methyl orange Merck 8 .O 105 4-chlorophenol Degussa P25 0.5 140 Methylene blue Degussa P25 0.15 132

Malonic acid Degussa P25 0.8 96 2-chlorophenol CERAC 2.0 114

Cyanides Home prepared 2.0 113 4-nitrophenol Degussa P25 2.0 56

Dimethylphenols Degussa P25 1 .o 112

When Ti02 catalyst is immobilized on supports, there exists an optimum thickness

for the catalyst film. There are two types of catalyst illumination [102]. In the first

arrangement (SC illumination), UV light is introduced from the catalyst support

(substrate) side, while in the second arrangement light is introduced from the liquid

side (LC illumination). In SC illumination, catalyst support must be optically

transparent, such as quartz and Pyrex glass. The advantage of this arrangement is that

UV light doesn't pass through the reaction solution before reaching the catalyst film,

530


and therefore, light intensity will not be attenuated due to the absorption by reaction

medium. Influence of catalyst film thickness, h, on the photocatalytic reaction can be

described by following equation [ 1031:

... (33)

and optimum catalyst layer thickness is given by

... (34) wffp De 1 k, )

k , ffp-- De

hop, =

In the LC illumination case, the effect of catalyst film thickness on the degradation

rate is given by [ 1031:

Obviously photocatalytic reaction rate reaches a saturation value as the catalyst

layer thickness increases. However, if the film is too thick, the strength of catalyst

adhesion is poor and is likely be detached from its supports. Catalyst film with less

than 6pm thickness was widely used in laboratory studies [117-1211, because

experimental result indicated that 95% of the incident light intensity has been

absorbed by the catalyst film with thickness of 4 . 8 ~ of P25 Ti02 [ 1031.

Dissolved oxygen

According to the principles of photocatalytic reaction, the rate of oxidation by holes

has to be balanced by the reduction rate of the photogenerated electrons. Therefore,

either the oxidation or reduction reaction can be the rate-governing step. Accumulated

electrons may also result in increasing rate of electron-hole recombination step,

thereby, decreasing in the hole-initiated oxidation of organic molecules. Thus, the

excess in photogenerated electron must be removed in order to avoid recombination.

Hence, electron scavenger (usually dissolved oxygen) is necessary. It was found [77]

531


that photocatalytic activity nearly completely suppressed in the absence of dissolved

oxygen, and the steady-state concentration of dissolved oxygen had a profound effect

on the rate of photocatalyzed decomposition of organic compounds. Alberici and

Jardim [ 1221 found that decomposition of phenol in non-aerated solutions containing

Ti02 was much slower compared with aerated ones. Sabate et al. [lo61 did not

observe photocatalytic degradation of 3-chlorosalicylic acid when pure N2 was

bubbled through the solution. Hsiao et al. [29] found that the observed rate of

disappearance of dichloro- and trichloromethane was much faster when the solution

was well purged with oxygen. Sclafani et al. [83] found that the concentration decay

of phenol was much faster when oxygen was bubbled in the solution containing

anatase Ti02 instead of He. Similarly, Augugliaro et al. [62, 1161, found that by

increasing the partial pressure of oxygen in the He-02 mixture bubbled through a

phenol solution containing T i 4 catalyst, the reaction rate constant increased and

remained constant for poZ> 0.6. It is important, therefore, to provide sufficient oxygen

in the organic solution containing the Ti02 catalyst to prevent accumulation of

electrons on the TiOz surface. The major role of oxygen in photocatalytic degradation

of organic compounds is to act as a scavenger for continuous electron removal from

photocatalysts [ 1231.

Gerischer and Heller [81] studied in detail the role of oxygen in the photo

destruction of organics on catalyst surfaces. They developed kinetic models to predict

the maximum electron uptake by oxygen and found that the latter depends on catalyst

particle sizes and oxygen concentration in the solution. Early studies have shown that

the adsorption of oxygen on illuminated Ti02 surface depends on the number of

hydroxyl groups of the surface [124-1261. The dependence of degradation rate

constants of organics on the dissolved oxygen concentration can be well described by

the Langmuir-Hinshelwood equation (Table 6 summarizes the K02 values reported in

literature):

However, in photocatalytic removal of dissolved metal ions, dissolved oxygen is a

strong competitor for metal ions to scavenge the photogenerated electrons. So, the

532


presence of dissolved oxygen may significantly inhibit the photoreduction of metal

ions [37,47, 127, 1281.

Table 6. Adsorption constant of dissolved oxygen on TiOl catalyst.

Pollutant Catalyst KO, (atm-') Ref.

4-CP Degussa P25 10.5 93

4-CP" Degussa P25 4.4 140

Phenol Wako Chem. 11.1 61

Methyl orange Merck 4.2 105

4-CP Degussa P25 17.6 99

4-NP Degussa P25 9.98 99

Phenol Degussa P25 12.7 99

UV Light Intensity

The incident light intensity determines the photogenerated electrodhole pair

concentration, and thus the hydroxyl radical formation rate. Therefore, at sufficiently

low levels of illumination (catalyst dependent), degradation rate is of first-order in

intensity. At high intensity level, on the other hand, the reaction rate increases with

the square root of intensity because of the increasing electrodhole pair recombination

(Equation 7) during their migration to the particle surface and recombination of

hydroxyl radicals resulting from their high concentrations.

OH' +OH' + H202 . . . (37)

As a consequence, this is obviously detrimental to the photocatalytic process as it

results in the decrease in quantum efficiency. The optimal light power utilisation

should be in the domain where degradation rate is proportional to the incident light

intensity. Several previous studies [35, 50, 941 have demonstrated above conclusion.

The transition points between these regimes, however, will vary with the

photosystem. In the degradation of 3-CP in TiOz aqueous suspensions, the initial

degradation rate was reported to be proportional to radiant flux for values smaller than

20 mW/cm2, while above this value it was proportional to its square root [94]. The

initial degradation rate of 4-CP against radiant flux was also investigated in the range

533


of 2-50 mW/cm2 [50], where similar rate shift from first-order to half-order in

intensity was observed.

Table 7. Photocatalytic activation energies (E) reported in literature.

Reactant Catalyst E (kJ/mol) T (“1 Ref. 4-CP Fixed 20.6 10-60 119 4-CP Slurry 16 15-55 18 4-CP Slurry 5.5 10-45 50 4-CP Slurry 13.7 15-50 99

Methyl orange Slurry 18 27-46 105 Malonic acid Slurry 9.99 21-51 96

phenol Slurry 10 20-50 80 phenol Slurry 11.8 15-50 99

Salicylic acid Slurry 1 1 25-45 95

Oxalic acid Slurry 13 5-70 55 Formic acid Slurry 16.7 150

xylenols Slurry 8.8 6-60 112 4-Np Slurry 1.42 15-50 99

Benzoic acid Slurry 9.13 15-50 59 SBS Slurry 13.4 25-38 110

Methylene blue Fixed 60.6 16-52 118

2-propanol Slurry 31 4-40 67

Reaction Temperature

In photocatalysis, irradiation is the primary source of electron-hole pair generation at

ambient temperature as the band-gap energy is too high to be overcome by thermal

activation. The true activation energy should be very small or even zero. Thus, the

overall photocatalytic reaction (Equation 2) is usually found to be temperature

insensitive. Increasing the reaction temperature may increase the oxidation rate of

organic compounds at the interface. The dependence of degradation rate on

temperature is reflected by the magnitude of activation energy. Table 7 lists the

apparent activation energies of some organic compounds in reaction photocatalyzed

by TiOz. All the reported activation energies [less than 21 kT/mol] are much lower

than those of ordinary thermal reactions except the degradation of methylene blue

reported by Naskar et al. [118]. These values are quite close to that for a hydroxyl

radical reaction [95], suggesting that the photodegradation of most organic pollutants

534


be governed by hydroxyl radical reaction. This indicates that the temperature

influence is quite weak.

This is because that the low thermal energy (kT = 0.026 eV at room temperature)

has almost no contribution to the activation of Ti02 due to its high bandgap energy

(3.2 eV). Pichat and Herrmann [129] reported that, at ‘high’ temperature (T > 343K),

even a negative apparent activation energy was found in photocatalytic degradation.

A possible explanation is that the adsorption of the reactants may have a significant

effect on the degradation rate because adsorption capacity is usually affected

inversely by increasing temperature. As a consequence, contrary to thermal reactions

there is no need to heat the system. This absence of heating is very attractive for

photocatalytic processes carried out in aqueous media, especially for photocatalytic

water purification, because there is no need of additional energy in heating large

volume of water, which has a high value of heat capacity.

p H in Solution and Anions

The rate of photocatalytic degradation of an organic compound not only depends on

the parameters discussed above, but also on pH and presence of interfering adsorbing

species. Theoretically, pH value of the solution has strong influence on all oxide

semiconductors, including the surface charge on the solid catalyst particles, the size of

the aggregates formed and the band-gap energies of the conductance and valence

bands. Besides, pH also influences the adsorption properties of organic compounds

and their dissociating state in solution. In literature, influence of pH on the

photocatalytic degradation rate is diversified. Higher reaction rates for various

photocatalytic processes have been reported at both low and high pH [18-20] and no

general conclusions have been obtained till now. Typically, reaction rate varied by

less than one order of magnitude from one end of the pH range to the other.

Maximum degradation rate of 4-NP was reported [I301 at pH 4.5, but the degradation

rate increased insignificantly with the increase of pH value beyond 10. Augugliaro et

a]. [ 1313 investigated the effect of pH in the range of 3 to 11 on the degradation of 4-

NP in Ti02 suspension and found that the degradation rate decreased slightly with

increasing pH value. Herrmann et al. [55 ] reported that both low and high pH were

not favorable to the photocatalytic oxidation of oxalic acid in the TiO, aqueous

535


suspensions. Kim and Anderson [ 1321 found the best pH value for both photocatalytic

and photo-electrocatalytic degradation of formic acid over immobilized Ti02 film is

around its pK, value, i.e. pH = 3.75. They explained the result in terms of the

influence of pH on the adsorption of formate ion (HCOO') on Ti02 surface. Chen and

Ray [56] also found that both low and high pH were not suitable for the degradation

of 4-nitrophenol in P25 Ti02 suspensions.

The high O H ion content of the system enhances the electrodhole separation

through reaction (Equation 11). However, in alkaline system the photogenerated COz

will be trapped in the solution, and bicarbonate and carbonate are formed. According

to the carbonate equilibrium:

PKl = 6.3 H 2 C 0 3 ( C 0 2 -I- H 2 0 ) \ - H C O F + H + ...

PK1= 10.2 HCOT \ -Cot- + H' ... (39)

Between pH 6.3 and 10.2, the inorganic carbon is mainly present in the form HC03-

while pH >10.2 as CO:-. At pH < 6.3, molecular C02 dominates, which evaporates

easily in an open system. Hence, the reason for retardation of the photo-degradation at

pH > 6.3 may be an inhibiting effect of HCO</CO?, which are known as efficient

scavengers of hydroxyl radicals [58, 1331 through reactions (39) and (40) due to their

very high rate constants with the hydroxyl radicals (k = 3 . 9 ~ 1 0 ~ M ' i ' for carbonate

and k = 8 . 5 ~ 1 0 ~ M I S - ' for bicarbonate) [134]. On the other hand, they also can absorb

on the Ti02 surface and may undergo a competitive reaction (41) with the desired

oxidation of organic pollutants.

O H + HCOT 4 H 2 0 + C 0 ; - ... (40)

OH* +co:--oH-+co;- ... (41)

HC@- + h + - H C G ... (42)

Although, the photogenerated carbonated radical anion has shown to be an oxidant

itself [ 1351, its oxidation potential is mush less positive than that of the OH' radical

and is not able to play an important role in the photocatalytic processes. On the other

536


hand, according to Equation (11) in acidic system the effect of carbonate and

bicarbonate can be avoided, but the low OH' hinders the formation of hydroxyl

radicals and subsequently reduces the degradation rate. Another factor that may affect

the rate of photocatalytic reaction is the anions in the solution. Abdullah et al. [I361

have carried out one of the rigorous studies on the effect of anions on the rate of

photomineralization of several organic pollutants catalyzed by TiOz. Anions, reported

in literature, which have negative effect on the degradation rate, include sulfate,

chloride, phosphate, carbonate and bicarbonate ions [ 136-1391. Similar result was also

reported elsewhere [28, 29, 1161. The inhibition of above anions is probably through

either as hydroxyl radical scavengers or as permanent adsorbates which consequently

blocks the reaction sites.

Crystalline Structure and Particle Size

Although TiOl is commonly used in the photocatalytic processes, the choice of which

crystalline form is still important. TiOz has three crystalline structures, anatase, rutile

and brookite, although brookite is not commonly available. Several studies have noted

that rutile is not an active catalyst [61, 621, although it is not clear why this should be.

This may be due, in part, to the small bandgap energy difference. The bandgap energy

for rutile is 3.0 eV, as compared to 3.2 eV for anatase. Thus the oxidatiodreduction

potentials are slightly less for the rutile phase and, thermodynamically rutile form

may not favor some reactions. MiIls and Morris [I401 observed that the degradation

rate of 4-chlorohenol on mile TiOz is much lower than that on anatase Ti02 when the

crystalline phase was converted at high annealing temperature (>7OO0C).

Photocatalytic reaction is a process that is closely related with the adsorption of

species on the catalyst surface. Therefore, it is expected that the specific surface area

of the catalyst should have significant effect on its activity due to the increased area

for adsorption of organic substrate, HzO and OH', and the corresponding subsequent

generation of hydroxyl radicals. Meanwhile, catalyst particle size may also affect the

photocatalytic reaction rates through an effect on the probability of electronhole

recombination. As particle size increases, the mass-transfer distance before electron

or hole reacting at the surface of catalyst increases. Consequently, any decreased

degradation rate for large particles may result from the higher degree of electron-hole

537


recombination. However, this theory has not been found to be consistent. Chen [59]

compared the activities of two brands of TiO2, Degussa P25 (70% anatase, 30% rutile,

d, = 30 nm) and Hombikat WlOO (100% anatase, d, = 10 nm), under identical

conditions in photocatalytic degradation of phenol and its derivatives. Experimental

results demonstrated that P25 was more active that UV100. Therefore, besides the

crystalline, surface area and particle size, there may be other controlling parameters

(such as the methods of synthesis) that must be considered when evaluating different

brand Ti02 or Ti02 from different sources.

The Photocatalytic Reactor In spite of the potential of this technology, development of a practical water treatment

system has not yet been successfully achieved. Several factors impede the effective

design of photocatalytic reactor [ 1551. In photocatalytic reactors, it is necessary to

achieve efficient exposure of the catalyst to light irradiation. Without photons of

appropriate energy, the catalyst shows no activity. In fact, in a photocatalytic reactor,

besides conventional reactor complications such as mixing, mass transfer, reaction

kinetics, catalyst installation, etc., an additional engineering factor related to

illumination of catalyst becomes relevant. The problem of photon energy absorption

has to be considered regardless of reaction kinetics mechanisms. The high degree of

interaction among the transport processes, reaction kinetics, and light absorption leads

to a strong coupling of physicochemical phenomena and a major obstacle in the

development of photocatalytic reactors. The illumination factor is of utmost

importance since the amount of catalyst that can be activated determines the water

treatment capacity of the reactor [156].

One major barrier to the development of a photocatalytic reactor is that the

reaction rate is usually slow compared to conventional chemical reaction rates, due to

low concentration levels of the pollutants [157]. Other crucial hurdle is the need to

provide large amounts of active catalyst inside the reactor. Even though the effective

surface area of the porous catalyst coating may be high, there can only be a thin

coating (about 1 pm thick) applied to a surface. Larger thickness of catalyst layer

washes away during experiments due to poor adhesion. Thus, the amount of active

catalyst in the reactor is limited and, even if individual degradation processes can be

538


made relatively efficient, the overall conversion efficiency will still be low [ 1581.

This problem severely restricts the processing capacity of the reactor and the time

required to achieve high conversions are measured in hours, if not days [ 1591.

In numerous investigations, an aqueous suspension of the catalyst particles in

immersion or annular type photoreactors has been used. However, the use of

suspensions requires the separation and recycling of the ultrafine catalyst from the

treated liquid that is usually an inconvenient, time consuming expensive process. In

addition, the depth of penetration of UV light is limited because of strong absorption

by catalyst and dissolved pollutants. One solution to the above problem is to

immobilise the catalyst onto a fixed transparent support. The immobilisation of

catalyst, however, generates another problem. The reaction occurs at the liquid-solid

interface and the overall rate may be limited to mass transport of the pollutant to the

catalyst surface [ 1561.

In view of the above, new reactor configurations must address two most important

parameters: (i) light distribution inside the reactor through absorbing and scattering

liquid to the catalyst, and, (ii) providing high surface areas of catalyst coating per unit

volume of reactor. The new reactor design concepts must provide a high ratio of

activated immobilised catalyst to illuminated surface and also must have a high

density of active catalyst in contact with liquid to be treated inside the reactor. A

number of photocatalytic reactors have been patented in recent years, but only few

[ 153, 1541 so far been developed to pilot scale level. Based on the arrangement of the

light source and reactor vessel, all these reactor configuration fall under the categories

of immersion type with lamp(s) immersed within the reactor [160, 1611, external type

with lamps outside the reactor [or distributive type with the light distributed from the

source to the reactor by optical means such as reflectors or optical fibres [ 162- 1631.

Process Efficiency and Catalyst Reusability Photocatalytic degradation of numerous organic species has been investigated in

laboratory scale. Obviously, it is necessary to compare the efficiencies of different

processes. Attempts have been made to devise measurable parameters that may be

helpful to compare results between research groups and those derived from other

wastewater treatment technologies. Bolton and co-workers [ 141, 1421 proposed that

539

D. Chen, M . Sivakumar and A. K . Ray

“electrical energy per unit mass (EEM)” and “electrical energy per order (EWO)”

could be used as appropriate figures of merit for zero-order and first-order kinetics,

respectively. The former was defined as the electrical energy (kWh) needed to

degraded a kilogram of pollutants, while the latter was the electrical energy (kWh)

needed to degrade pollutants in 1 m3 aqueous solution by one order of magnitude. The

above figures of merit may help to compare the efficiency of the entire photocatalytic

process. However, they do not provide a direct indication of the efficiency of an ultra-

bandgap photon to drive the semiconductor-catalyzed process of reaction (Equation

2). To do this, quantum yield is extensively adopted to evaluate the process efficiency

in photocatalytic processes.

The quantum yield, 9, for a photocatalytic reaction is defined as the number of

molecules of target compound converted divided by the number of photons of light

absorbed by photocatalyst [ 1431:

3. rate of reaction organics, (mol I m s) - ... (43) N

- absorption rate of radiation cp=

photon, (moll m 3 .s)

Ferrioxalate {[Fe(C204)3]3-; for UV and visible region to 500nm) and Reineck’s salt

{ [Cr(NH&(SCN)J; for visible region} [ 151 have been used as standards to measure

the photon flow incident on the inner front window of a photolytic cell. These are

known as chemical actinometers because the quantum yield from these substances is

rather insensitive to the changes in temperature, reactant concentration, photon flow

and wavelength of the incident light. Procedures are well established, analysis is

simple, and the precision in quantum yield data is unquestioned [143]. Normally, the

maximum quantum yield possible is unity. However, if the photocatalytic reaction

initiates a chain reaction, the quantum yield can be considerably great than unity.

In heterogeneous photocatalysis, quantum yield can be described in the same

manner as for homogeneous photocatalysis if the number of actual absorbed photons

by the solid photocatalyst can be assessed by some spectroscopic means. Practically,

however, usage of the term quantum yield as defined in above equation encounters

many difficulties since the number of absorbed photons, Nphomn, is impossible to

determine experimentally due to reflection, scattering, transmission and absorption by

540


the suspended particulate. A metal oxide material such as Ti02 particulate can never

absorb all the incident photon flux from a given source [144]. It was reported that

only about 60-65% of the incident photons could be absorbed by the Ti02 in aqueous

solution [145]. In order to circumvent the inherent difficulties encountered in the

precise evaluation of the number of quanta absorbed by the photocatalyst, Serpone

and co-workers [ 144-1481 proposed the concept of relative photonic efficiencies.

They chose the initial photo conversion of phenol as standard process, and Degussa

P25 Ti02 as standard photocatalyst; and defined the relative photonic efficiency as:

initial disapperance rate of substrate initial disappearance rate of phenol

5 * = . . .. . (44)

where both initial rate are obtained under exactly identical conditions. The, 6, , value

can be converted into photochemically defined quantum yield, (Psubsme, once the

quantum yield of phenol, (Pphenol, on Degussa P25 Ti02 can be determined:

... (45) - psubsmie - 6 r q p h e n o i

The use of relative photonic efficiency, 5, , renders comparison of process

efficiencies (relative to phenol) between studies carried out in different laboratories

possible and can determine which organic substrate is transformed most efficiently.

However, all experiments must be conducted under exactly identical conditions, and

the results did not take into account of the effect of intermediates on the degradation

of their parent substrates because the initial rates are used to evaluate the relative

photonic efficiencies. In practical, intermediates formed during photoreaction may

hinder the degradation of their parent compounds and their distribution may depend

on experimental conditions. Photocatalytic process efficiency was also reported in

terms of the apparent quantum yield, which was defined by:

2 ... (46) - ri -- disappearance rate of substrate (mourn s)

rate of photon incident on window (einsteidm s) (Pappr = 2 Ftoml

When a polychromatic light source is used, the denominator in this equation can be

written as:

541


where, N, is the Avogadro number (6.023 x cv is the light rate (3x10' d s ) , h is

the Planck's constant ( 6 . 6 3 ~ 1 0 ~ ~ J s), h is the wavelength (m), and W(h) the light

intensity at wavelength h arriving at the reactor window.

All above evaluation parameters provide some idea of the efficiency of

photocatalytic process. However, it must be pointed out that, in heterogeneous

photocatalysis, reaction rate is influenced by all the parameters. As a result, even for

the same reaction system, the process efficiency reported by one research group might

be quite different from thqse obtained by others due to the difference in experimental

conditions. Nevertheless, photocatalytic process efficiency reported in literature in

any one of above parameters is very low, less than 8% [95, 106, 109, 149, 1501. This

is one of the main hindrances to the commercialization of photocatalysis. Therefore,

improving process efficiency is one of the most important challenges in this field.

When Ti02 is immobilized onto catalyst supports, the process efficiency is much

lower. The immobilized Ti02 catalyst was usually 2 to 3 times less effective than the

freely suspended one [35, 118, 151, 1521. It is more likely that this discrepancy

resulted from the low availability of active sites, and mass transfer resistance in

immobilized catalyst system. However, it can be concluded that even if suspensions

remain more efficient, the photo-activity of immobilized Ti02 samples was of the

same order of magnitude. From an application point of view, the interest of having

stable, supported photocatalysts that avoid filtration may actually compensate the

drawback of a lower activity.

Reusability of photocatalyst is very important. Phenol is an ideal standard organic

compound for testing sustained activity of the catalyst since during the photocatalytic

reaction of phenol the pH value may almost be kept constant, and no anions are

accumulated. As a result, the effect of pH and anion on the photocatalytic activity can

be avoided. This renders the comparison to be more reliable. We [59] found that

activity of P25 Ti02 was quite stable because the decrease of activity in five

successive runs (about 28 hours of illumination) of the degradation of phenol was

within experimental error. Al-Sayyed et al. [50] tested sustained activity of Degussa

542


P25 catalyst in ten successive cycles (about 20 hours of successive illumination) using

4-CP as model compound. Small decline in photocatalytic activity was observed in

their investigation. But, they attributed this decline to the inhibiting effect of C r ions,

whose concentration gradually increased, and to some Ti02 loss owing to the

samplings for analyses.

Summary and Outlook Heterogeneous photocatalysis is a green technology and may not require the addition

of consumable reagents. It requires a recoverable, environmentally benign solid Ti02

catalyst and can also be activated by solar photons. The process involves electron-

hole pair generation initiated by band-gap excitation. The photogenerated holes and

electrons have high oxidation and reduction potentials respectively, and therefore, can

be extensively applied for environmental detoxification. It offers complete destruction

of almost all-organic pollutants in water and air, including elimination of many

inorganic compounds as well. The oxidation andor reduction technologies at times

are the only method available to achieve complete destruction of the contaminants,

particularly at low concentration. The photogenerated holes may react with the

adsorbed hydroxyl groups or water to form hydroxyl radicals, which subsequently

mineralize organic species. On the other hand, photogenerated electrons are usually

trapped by dissolved oxygen, or by other absorbed electron-acceptors (such as metal

ions). In the latter case, toxic metal ions can be removed. Based on the above theory,

waste streams containing organic species and metal ions can be detoxified making

photocatalysis an important process technology.

It is technically feasible to remove or mineralize a wide range of organic and

inorganic compounds from contaminated soil, water and air using photocatalytic

process. Hence, photocatalytic process has the potential to be a commercially viable

detoxification technology. This is very exhilarating. However, only a few applications

have been reported [153, 1541. Clearly the photocatalytic processes are still in the

early stages of development and implementation. In order for this technology to be

commercialized successfully, the photo-efficiency of the process has to be improved

significantly. Although Degussa P25 TiOl is the most efficient photocatalyst found to

date, its activity is still much lower than the requirement in large-scale applications.

543


Hence, modification of the catalyst andor development of other catalysts is one of the

keys for the future success of the process. Actually, a great efforts being made to

improve the catalyst performance including annealing by heat treatment, doping with

metal ions, alternative preparation methods, deposition of noble metals or metal oxide

on the Ti02 surface, and other surface modifications. But the improvement achieved

till now is not significant. Prior to its commercialization, a desired photocatalytic

reactor configuration must also be designed in which light can be distributed

uniformly to a large surface area of catalyst. Besides, TiOz absorbs radiation below

385 nm. Unfortunately, this represents only 3% of the solar spectrum at the Earth’s

surface. If Ti02 (or other) catalyst could be modified to absorb at longer wavelengths

to allow use of sunlight as light source, a considerable reduction in the operation cost

can be realized. All these would make photocatalytic process much more attractive

than conventional waste treatment technologies.

References 1.

2.

3.

4.

5 .

6.

7.

8.

9.

10.

11.

12.

13.

14. 15.

544

~~

Conner, J. R. 1990. Chemical fixation and solidification of hazardous wastes, Van Nostrand Reinhold, New York. Matthews, R. W. 1992. Photocatalytic oxidation of organic contaminants in water: An aid to environmental preservation. Pure and Appl. Chem., 64, pp1285-1290. Mills, A. and S. L. Hunte. 1997. An overview of semiconductor photocatalysis. J. Photochem. Photobiol. A: Chem.. 108, ppl-35. Pelizzetti, E.; V. Maurino; C. Minero; V. Carlin; E. Pramauro; 0. Zerinati and M. L. Tosato. 1990 Photocatalytic degradation of atrazine and other striazine herbicides. Environ. Sci. Technol.. 24.

Shaub, W.M. and W. Tsang. 1983. Dioxin formation in incinerators. Environ. Sci. Technol., 17.

Legrini, 0.; E. Oliveros and A. M. Braun. 1993. Photochemical processes for water treatment. Chem. Rev., 93, pp671-698. Khan, S. R.; C. R. Huang and J. W. Bozelli. 1985. Oxidation of 2-chlorophenol using ozone and ultraviolet radiation. Environ. F’rog., 4, pp229-238. Schiavello, M. (Eds.) 1985. Photoelectrochemistry, Photocatalysis and Photo reactors, Fundamentals and Developments. Dordrecht, Reidel. Schiavello, M. (Eds.) 1988. Photocatalysis and Environment: Trends and Applications. Dordrecht. Kluwer. Ollis D. F.; E. Pelizzetti and N. Serpone, 1989. Heterogeneous photocatalysis in the environment: application to water purification. In Photocatalysis: Fundamentals and Applications, (Edited by Serpone N. and Pelivetti E.) Wiley Intersicence, New York, pp603-637. Serpone, N. and Pelizzetti, E. (Eds.)1989. Photocatalysis: Fundamentals and Applications. New York: Wiley. Ollis, D. F. and Al-Ekabi, H. (Eds.) 1993. Photocatalytic Purification and Treatment of Water and Air. Amsterdam: Elsevier. Pelizzetti, E. and Serpone, N. (Eds.) 1986. Homogeneous and Heterogeneous Photocatalysis. Dordrecht, Reidel. Gratzel. M., 1989. Heterogeneous Photochemical Electron Transfer. CRC press, Boca Raton, Fla. Braun, A. M.; M. T. Maurette and E. Oliveros. 1991. Photochemical Technology. Wiley, New York.

1559-1565

pp.721-730.


16.

17. 18.

19. 20.

21.

22.

23. 24.

25.

26. 27.

28.

29.

30.

31.

32.

33.

34.

35.

36.

37.

38.

39.

40.

41.

42.

Rajeshwar, K. and J. G. Ibanez. 1997. Environmental Electrochemistry, Fundamentals and Applications in Pollution Abatement. Academic Press, INC. Schiavello, M. (Eds.) 1997. Heterogeneous Photocatalysis. Chichester: John Wiley & Sons. Mills, A.; R. Davies and D. Woaley. 1993. Water purification by semiconductor photocatalysis. Chemical Society Reviews, 22, pp417-425. Fox, M. A. and Dulay, M. T., 1993. Heterogeneous photocatalysis. Chem. Rev., 93, pp341-357. Hoffmann, M. R.; S. T. Martin; W. Choi and D. W. Bahnemann. 1995. Environmental applications of semiconductor photocatalysis. Chem. Rev. 95, pp69-96. Hagfeldt, A. and Gratzel, M. 1995. Light induced redox reactions in nano-crystalline systems, Chem. Rev. 95.49. Howe, R. F. 1998. Recent developments in photocatalysis. Dev. Chem. Eng. Mineral Process., 6,

Ollis, D. F. 1985. Contaminant degradation in wafer. Environ. Sci. Technol., 19, pp480-484. Choi, W.. Termin, A. and M. R. Hoffmann. 1994. Role of metal ion dopants in quantum-sized TiOl. Correlation between photoreactivity and charge canier recombination dynamics. J. Phys. Chem., 98, pp13669-13679. Miller R. and Fox R. 1993. In Photocatalytic Purification and Treatment of Water and Air. Edited by D. F. Ollis, H. Al-Ekabi, pp573-578 Amsterdam: Elsevier. Fujishima, A. and K. Honda. 1972. Nature 37, p238. Pruden, A. L. and D. F. Ollis. 1983. Photoassisted heterogeneous catalysis: The degradation of trichloroethylene in water. J. Catal., 82, pp404-417. Pruden, A. L. and D. F. Ollis. 1983. Degradation of chloroform by photoassisted heterogeneous catalysis in dilute aqueous suspensions of TiO2. Environ. Sci. Technol., 17, pp628-631. Hsiao. C. Y.; C. L. k, and D. F. Ollis. 1983. Heterogeneous photocatalysis: Degradation of dilute solutions of dichloromethane. chloroform and carbon tetrachloride with illuminated Ti02 photocatalyst. J. Catal, 82, pp418-423. Ollis, D. F.; C. Hsiao; L. Budiman and C. L. Lee. 1984. Heterogeneous photoassisted catalysis: conversion of perchloroethyiene, dichloroethane, chloroacetic acids, and chlorobenzenes. J. Catalysis, 88, pp89-96. Turchi. C. S. and D. F. Ollis. 1989. Mixed reactant photocatalysis: Intermediates and mutual rate inhibition. J. Catal., 119, pp483-496. Turchi, C. S . and D. F. Ollis. 1990. Photocatalytic degradation of organic water contaminants: mechanisms involving hydroxyl radical attack. J. Catal. 122, pp178-192. Matthews. R. W. 1988. Kinetics of photocatalytic oxidation of organic solutes over titanium dioxide. J. C a d . 11 1, pp264-272. Matthews, R. W. 1988. An adsorption water purifier with in situ photocatalytic regeneration. J. Catal., 113, pp549-555. Matthews, R. W. and S. R. McEvoy. 1992. Photocatalytic degradation of phenol in the presence of near-UV illuminated titanium dioxide. J. Photochem. Photobiol. A: Chem., 64, pp231-246. Ollis D. F.; E. Pelimtti and N. Serpone. 1991. Destruction of water contaminants. Environ. Sci. Technol., 25, pp1523-1529. Prairie, M. R.; L. R. Evans; B. M. Stange and S. L. Martinez. 1993. An investigation of Ti02 photocatalysis for the treatment of water contaminated with metals and organic chemicals. Environ. Sci. Technol. 27, pp1776-1782. Hemnann, J. M.; J. Disdier and P. Pichat. 1988. Photocatalytic deposition of silver on powder titania: Consequences for the recovery of silver. J. Catal. 113, pp72-81. Fu, H., G. Lu and S. Li. 1998. Adsorption and photo-induced reduction of Cr(VD ion in Cr(VD-4- CP aqueous system in the presence of Ti02 as photocatalyst. J. Photochem. Photobiol. A: Chem.,

Khalil, L. B.; W. E. Moumd ans M, W. Rophael 1998. Photocatalytic reduction of environmental pollutant Cr (Vi) over some semiconductors under UVIvisible light illumination. Applied Catalysis B: Environmental, 17, pp267-273. Navio, J. A., Colon, G., Trillas, M., Peral, J., Domenech, X., Testa., J. J.. Padron., J.. Rodriguez, D and Litter, M. 1.. 1998. Heterogeneous photocatalytic reactions of nitrite oxidation and Cr (VI) reduction on iron-doped titania prepared by the wet impregnation method. Applied Catalysis B: Environmental, 16, pp187-196. Serpone, N.; Y. K. Ah-You; T. P. Tran; R. Harris; ,E. Pelizzetti and H. Hidaka. 1987. AM1 simulated sunlight photoreduction and elimination of Hg (11) and CH3Hg (11) chloride salts from aqueous suspensions of titanium dioxide. Sol. Energy, 39, pp491-498.

~ ~ 5 5 - 8 4 .

114, ~ ~ 8 1 - 8 8 .

545


43.

44.

45. 46. 47. 48.

49.

50.

51.

52.

53.

54.

55.

56.

57.

58.

59.

60.

61.

62.

63.

64.

65.

66.

67.

68.

Angelidis, T. N.; M. Koutlemani and I. Poulios 1998. Kinetic study of the photocatalytic recovery of Pt from aqueous solution by TiOl, in a closed-loop reactor. Applied Catalysis B: Environmental, 16, pp347-357. Kisch, H. 1989. What is photocatalysis? In Photocatalysis: Fundamentals and Applications. Edited by Nick Serpone, Ezio Pelizzetti. New York: Wiley ppl-7. Kisch, H. and H. Henning. 1983. EPA Newslett. 19, p23. Boer, K. W., 1990. Survey of Semiconductor Physics. Van Nostrand Reinhold, New York. pp249. Dingwang metal ions Barbeni, M.; E. Pramauro; E. Borgarello and N. Serpone. 1985. Photodegradation of pentachlorophenol catalyzed by semiconductor particles. Chernosphere. 14, pp 195-208. Augugliaro, V.; L. Palmisano; M. Schiavello; A. Sclafani; L. Marchese; G. Mama and F. Miano. 1991. Photocatalytic degradation of nitrophenols in aqueous titanium dioxide dispersion. Appl. Catal., 69, pp323-340. Al-Sayyed, G.; J. C. D’Oliveira and P. Pichat. 1991. Semiconductor-sensitized photodegradation of 4-cholophenol in water. J. Photochem. Photobiol. A: Chem., 58, pp99-114. Bekbolet, M.; M. Lindner; D. Weichgrebe and D. W. Bahnemann. 1996. Photocatalytic detoxification with the thin-film fixed-bed reactor (TFFBR): Clean-up of highly polluted landfill effluents using a novel Ti02-photocatalyst, Sol. Energy, 56, pp455-469. Ha, H. Y. and M. A. Anderson. 1996. Photocatalyhc degradation of formic acid via metal- supported titania. J. of Environmental Engineering, 122, pp217-221. Finklea, H. 0.. 1983. Photoelectrochemistry: Introductory concepts. J. Chem. Educ., 60, pp325- 327. Mills, A. and S. Moms. 1993. Photomineralization of 4-chlorophenol sensitized by titanium dioxide: A study of the initial kinetics of carbon dioxide photogeneration. J. Photochern. Photobiol. A: Chem., 71, pp75-83. Hemnann. J. M.; M. N. Mozzanega and P. Pichat. 1983. Oxidation of oxalic acid in aqueous suspensions of semiconductors illuminated with UV or visible light. J. Photochern. 22, pp333-343. Chen, D. W. and A. K. Ray. 1998. Photodegradation kinetics of 4-nitrophenol in Ti02 aqueous suspensions. Wat. Res., 32, p3223-3234. Turchi, C. S.; M. Mehos and J. Pacheco. 1993. In Photocatalytic purification and treatment of water and air. Ollis, D, F. and Al-Ekabi, H. Eds., Elsevier: Amsterdam, pp789-794. Mehos, M. S. and C. S. Turchi. 1993. Field testing solar photocatalytic detoxification on TCE- contaminated groundwater. Environ. Prog., 12, pp194-199. Chen, D. W. 1999. Development of photocatalytic process for wastewater treatment. Ph. D thesis, National University of Singapore. Krautler, B. and A. J. Bard. 1978. heterogeneous photocatalytic decomposition of saturated carboxylic acids on Ti02 powder. Decarboxylative route to alkenes. J. Am. Chem. SOC., 100,

Okamoto K., Yamamoto Y., Tanaka H. and Itaya A. 1985. Heterogeneous photocatalytic decomposition of phenol over Ti02 powder. Bull. Chem. SOC. Jpn. 58. pp2015-2022. Augugliaro, V. L. Palmisano; A. Sclafani; C. Minero and E. Pelizzetti. 1988. Photocatalytic degradation of phenol in aqueous titanium dioxide dispersions. Toxicol. Environ. Chem., 16, pp89- 109. Sclafani. A. and J. M. Hemnann, 1996. Comparison of the photoelectronic and photocatalytic activities of various anatase and rutile forms of titania in pure liquid organic phases and in aqueous solutions. J . Phys. Chem.. 100, pp13655-13661. Hemnann, J. M. 1995. Heterogeneous photocatalysis: an emerging discipline involving multiphase systems. Catalysis Today, 24, pp157-164. Bard, A. J. 1979. Photoelectrochemistry and heterogeneous photocatalysis at semiconductors. J. Photochemistry. 10, pp59-75. Palmisano, L. and A. Sclafani. 1997. Thermodynamics and kinetics for heterogeneous photocatalytic progresses. In Heterogeneous Photocatalysis. Edited by M. Schiavello, pp109-131. Chichester: John Wiley & Sons. Chlds, L. P. and D. F. Ollis. 1981. Photoassisted heterogeneous catalysis: Rate equations for oxidation of 2-methyl-2-butyl-alcohol and isobutane. J. Catal., 67, pp35-48. Alfano, 0. M.; M. I. Cabrera and A. E. Cassano. 1997. Photocatalytic reactions involving hydroxyl radical attack. 1: Reaction kinetics formulation with explicit photon adsorption effects. J. Catal.,

~~5985-5992.

172, ~~370-379 .

546


69.

70.

71.

72.

73.

74.

75.

16.

77.

78.

79.

80.

81.

82.

83.

84.

85.

86.

87.

88. 89.

90.

91.

Morishige, K.; F. Kanno; S. Ogawara and S. Sasaki. 1985. Hydrated surface of particulate titanium dioxide prepared by pyrolysis of &oxide. J. Phys. Chem., 89, pp4404-4408. Suda, Y. and T. Morimoto. 1987. Molecularly adsorbed Hz0 on the bare surface of Ti02 (rutile). Langmuir, 3, pp786-788. Jaeger, C. D. and A. J. Bard. 1979. Spin uapping and electron spin resonance detection of radical intermediates in the photodecomposition of water at Ti02 particulate systems. J. Phys. Chem., 83,

Fox, M. A. and C. Chen. 1981. Mechanistic features of the semiconductor photocatalyzed olifin-to- carbonyl oxidative cleavage. J. Am. Chem. SOC., 103, pp6757-6759. Matthews, R.W. 1984. Hydroxylation reactions induced by near-ultraviolet photocatalysis of aqueous titanium dioxide suspensions. J. Chem. SOC., Faraday Trans. 1,80, pp457-471. Al-Ekabi, H.; N. Serpone; E. Pelizzetri; C. Minero; M. A. Fox and R. 8. Draper. 1989. Kinetic studies in heterogeneous photocatalysis: 2. Ti02-mediated degradation of 4-chlorophenol alone and in a three-component mixture of 4-chlorophenol, 2,4-dichlorophenol, and 2,4,5-trichloropheno1 in air-equilibrated aqueous media. Langmuir, 5, pp250-255. Matthews, R. W., 1990. Purification of water with near-UV illuminated suspensions of titanium dioxide. Wat. Res., 24, pp653-660. Low, G. K. C., S. R. McEvoy and R. W. Matthews. 1991. Formation of nitrate and ammonium ions in titanium dioxide mediated photocatalytic degradation of organic compounds containing nitrogen atoms. Environ. Sci. Technol., 25, pp460-467. Wang, C. M., A. Heller and H. Gerischer. 1992 Palladium catalysis of 02 reduction by electrons accumulated on Ti02 panicles during photoassisted oxidation of organic compounds. J. Am. Chem. SOC., 114, ppS230-5234. Al-Ekabi, H. and N. Serpone. 1988. Kinetic studies in heterogeneous photocatalysis. 1. Photocatalytic degradation of chlorinated phenols in aerated aqueous solutions over Ti02 supported on a glass matrix. J. Phys. Chem., 92, pp5726-5731. Augugliaro, V.; E. Davi, L. Palmisano; M. Schiavello and A. Sclafani. 1990. Influence of hydrogen peroxide on the kinetics of phenol degradation in aqueous titanium dioxide suspension.

Okamoto K.; Y. Yamamoto; H. Tanaka and A. Itaya. 1985. Kinetics of heterogeneous photocatalytic decomposition of phenol over anatase Ti02 powder. Bull. Chem. SOC. Jpn. 58,

Gerischer, H. and A. Heller. 1991. The role of oxygen in photooxidation of organic molecules on semiconductor particles. J. Phys. Chem., 95, pp5261-5267. Tanaka, K.; T. Hisanaga and K. Harada. 1989. Photocatalytic degradation of organohalide compounds in semiconductor suspension with added hydrogen peroxide. New J . Chem., 13, pp5-7. Sclafani, A.; L. Palmisano and E. Davi. 1990. Photocatalytic degradation of phenol by Ti02 aqueous dispersions: rutile and anatase activity. New J. Chern., 14, pp265-268. Wei. T. Y.; Y. Y. Wang and C. C. Wan. 1990. Photocatalytic oxidation of phenol in the presence of hydrogen peroxide and titanium dioxide powders. J. Photochem. Photobiol. A: Chem., 55,

Matthews, R. W., 1991. Photooxidative degradation of coloured organics in water using supported catalysts Ti02 on sand. Wat. Res., 25, pp1169-1176. Martyanov, I. N.; E. N. Savinov and V. N. Parmon. 1997. A comparative study of efficiency of photooxidation of organic contaminants in water solutions in various photochemical and photocatalytic systems. 1. Phenol photooxidation promoted by hydrogen peroxide in a flow reactor. J. Photochem. Photobiol. A: Chem., 107. pp227-231. Ireland, J. C. and J. Valinieks. 1992. Rapid measurement of aqueous hydroxyl radical concentrations in steady state OH' flux systems. Chemosphere, 25, pp383-396. Cox, R. A,; R. G. Dement and M. R. Williams. 1980. Environ. Sci. Technol., 14, pp57. Baheni. M.; C. Minero; E. Pelizzetti; E. Borgarello and N. Serpone. 1987. Chemosphere, 16, pp2225. Fox, M. A., 1993. The role of hydroxyl radicals in the photocatalyzed detoxification of organic pollutants: Pulse radiolysis and time-resolved diffuse reflectance measurements. In Photocatalytic Purification and Treatment of Water and Air. Edited by D. F. Ollis, H. AI-Ekabi, pp163-167. Amsterdam: Elsevier. Minero, C., F. Catozzo and E. Pelizzetti. 1992. Role of adsorption in photocatalyzed reactions of organic molecules in aqueous Ti02 suspensions. Langmuir, 8, pp481-486.

~~3146-3152.

Appl. Catal., 65, ~~101-116 .

~~2023-2028.

~ ~ 1 1 5 - 1 2 6 .

547


92.

93.

94.

95.

96.

97.

98.

99.

100.

101.

102.

103.

104.

105.

106.

107. 108.

109.

110.

111.

112.

113.

114.

115.

116.

Foster, N. S.; R. D. Noble and C. A. Koval. 1993. Reversible photoreductive deposition and oxidative dissolution of copper ions in titanium dioxide aqueous suspensions. Environ. Sci.

Foster, N. S.; R. D. Noble and C. A. Koval. 1993. In Photocatalytic Purification and Treatment of Water and Air. Edited by D. F. Ollis, H. Al-Ekabi, pp365-373. Amsterdam: Elsevier. D’Oliveira J. C.; G. Al-Sayyed and P. Pichat. 1990. Photodegradation of 2- and 3-chlorophenol in Ti02 aqueous suspensions. Environ. Sci. Technol. 24, pp990-996. Matthews, R. W., 1987. Photooxidation of organic impurities in water using thin films of titanium dioxide. J. Phys. Chem., 91, pp3328-3333. Inel Y. and A. Okte. 1996. Photocatalytic degradation of malonic acid in aqueous suspensions of TiOz: an initial kinetic investigation of C02 photogeneration. J. Photochem. Photobiol. A: Chem.,

Matthews R. W. 1986. Photo-oxidation of organic material in aqueous suspensions of titanium dioxide. Wat. Res., 20. pp569-578. Chen, H. Y; 0. zahraa, M. Bouchy; F. Thomas and J. Y. Bottero. 1995. Adsorption properties of Ti02 related to the photocatalytic degradation of organic contaminants in water. J. Photochem. Photobiol. A: Chem., 85, pp179-186. Chen, D. W. and A. K. Ray, 1999. Photocatalytic kinetics of phenol and its derivatives over UV irradiated TiO2. Appl. Catalysis B: Environmental, 23, p143-157. Turchi, C. S. and Ollis D. F. 1988. Photocatalytic reactor design: an example of mass-transfer limitations with an immobilized catalyst. J. Phys. Chem., 92, pp6852-6854. Sclafani, A. Brucato, A and Rizzuti L. 1993. Mass Transfer Limitations in a Packed Bed Photoreactor Used for Phenol Removel. In Photocatalytic Purification and Treatment of Water and Air; Ollis, D. F. and Al-Ekabi, H. Eds.; Elsevier: Amsterdam, pp533-545 Ray, A. K. and Antonie A. C. M. Beenackers, 1997. Novel swirl-flow reactor for kinetic studies of semiconductor photocatalysis. AIChE J.. 43, pp2571-2578. Chen, D. W., Li, F. M. and A. K. Ray, 2000. Effect of mass transfer and catalyst layer thickness on photocatalytic reaction, AIChE, 46.5. Theurich, J.; M. Lindner and D. W. Bahnemann. 1996. Photocatalytic degradation of 4- chlorophenol in aerated aqueous titanium dioxide suspensions: a kinetic and mechanistic study. Langmuir, 12, pp6368-6376. Chen, L. C. and T. C. Chou. 1993. Kinetics of photodecolorization of methyl orange using titanium dioxide as catalyst. Ind. Eng. Ckm. Res., 32, pp1520-1527. Sabate, J.; M. A. Anderson; H. Kikkawa; M. Edwards and C. G. Hill. 1991. A kinetic study of the photocatalytic oxidation of 3-chlorosalicylic acid over Ti02 membrane supported on glass. J. Catal., 127, pp167-177. Satterfield, C. N. 1980. Heterogeneous Catalysis In Practice. McGraw-Hill: New York. Matthews, R. W. 1986. Photocatalytic oxidation of chlorobenzene in aqueous suspensions of titanium dioxide. J. Catal. 97, pp565-568. Lu, M. C.; G. D. Roam; J. N. Chen; and C. P. Huang. 1993. Factors affecting the photocatalytic degradation of dichlorvos over titanium dioxide supported on glass. I. Photochem. Photobiol. A: Chem., 76, pp103-110. Sangchakr, B.; T. Hisanaga. And K. Tanaka. 1995. Photocatalytic degradation of sulfonated aromatics in aqueous Ti02 suspension J. Photochem. Photobiol. A: Chem., 85, pp187-190. Vidal, A.; J. Herren; M. Romero; B. Sanchez and M. Sanchez. 1994. Heterogeneous photocatalysis: Degradation of ethylbenzene in Ti& aqueous suspensions. J. Photochem. Photobiol. A: Chern., 79, pp213-219. Terzian, R. and N. Serpone. 1995. Heterogeneous photocatalyzed oxidation of creosote components: Mineralization of xylenols by illuminated Ti02 in oxygenated aqueous media. J. Photochern. Photobiol. A: Chem., 89, pp163-175. Augugliaro, V., Loddo, V. Marci, G., Palmisano, L. and Lopez-Munoz, M. J. 1997. Photocatalytic oxidation of cyanides in aqueous titanium dioxide suspensions. J. of Catal., 166, pp272-283. Ku, Y.; R. M. Leu and K. C. Lee.. 1996. Decomposition of 2-chlorophenol in aqueous solution by UV irradiation with the presence of titanium dioxide. Wat. Res., 30, pp2569-2578. Wei, T.Y. and C. C. Wan. 1991. Heterogeneous photocatalytic oxidation of phenol with titanium dioxide powders. Ind. Eng. Chem. Res., 30, pp1293-1300. Augugliaro, V.; L. Palmisano and M. Schiavello. 1991. Photon absorption by aqueous Ti02 dispersion contained in a stirred photoreactor. AIChE J., 37, pp1096-1100.

Te~hnol., 27, pp350-356.

96, ~ ~ 1 7 5 - 1 8 0 .

548


117.

118.

119.

120.

121.

122.

123.

124.

125.

126.

127.

128.

129.

130.

131.

132.

133.

134.

135.

136.

137.

138.

139.

Lu, M. C.: G. D. Roam; J. N. Chen and C. P. Huang. 1995. Photocatalytic mineralization of toxic chemicals with illuminated TiOz. Chem. Eng. Corn . , 139, ppl-13. Naskar, S ; S. A. Pillay and M. Chanda. 1998. Photocatalytic degradation of organic dyes in aqueous solution with Ti02 nanoparticles immobilized on foamed polyethylene sheet. J. Photochem. Photobiol. A: Chem., 113. pp257-264. Hofstadler K.; R. Bauer; S. Novalic and G. Heisler. 1994. New reactor design for photocatalytic wastewater treatment with Ti02 immobilized on fused-silica glass fibers: photomineralization of 4- chlorophenol. Environ. Sci. Technol., 28, pp670-674. Nogueira, F. P. R. and W. F. Jardim. 1996. Ti02-fixed-bed reactor for water decontamination using solar energy. Sol. Energy, 56, pp471-477. Femandez, A.; G. hsalet ta and V. M. Jimenez. 1995 Preparation and characterization of Ti02 photocatalysts supported on various rigd supports (glass, quartz and stainless steel). Comparative studies of photocatalytic activity in water purification. Applied Catalysis B: Environmental, 7,

Alberici, R. M. and W. F. Jardim. 1994. Photocatalytic degradation of phenol and chlorinated phenols using Ag-Ti02 in a slurry reactor. Wat. Res., 28, pp1845-1849. Sukharev, V. and R. Kershaw. 19%. Concerning the role of oxygen in photocatalytic decomposition of salicylic acid in water. J. Photochem. Photobiol. A: Chem., 98, pp165-169. BicWey, R. I. and F. S . Stone. 1973. Photoadsorption and photocatalysis at rutile surface: I, photoadsorption of oxygen. J. Catal., 31, pp389-397. Boonstra. A. H., and Mutsaers. 1975 Relation between the photoadsorption of oxygen and the number of hydroxyl groups on a tatanium dioxide surface. 1. Am. Chem. SOC., 103, pp6324-6329. Munuera, G., V. Rives-Amau and A. Saucedo 1979. Photo-adsorption and photo-desorption of oxygen on highly hydroxylated Ti02 surfaces. J. Chem. SOC. Faraday Trans. I, 75, pp736-747. Munoz, J. and X. Domenech 1990. Ti02 catalyzed reduction of Cr(VI) in aqueous-solutions under ultraviolet illumination. J. Appl. Electrochem., 20, pp518-521. Lin, W.-Y.; C. Wei and K. Rajeshwar 1993. Photocatalytic reduction and immobilizgion of hexavalent chromium at titanium dioxide in aqueous basic-media. J. Electrochem. SOC., 140,

Pichat, P. and J. M. Hemnann. 1989. In Photocatalysis, Fundamentals and Applications, Ed. Serpone. N. and Peliuetti, Wiley, New Yo&, pp217. Palmisano, L; V. Augugliaro; M. Schiavello and A. Sclafani. 1989. Influence of acid-base properties on photocatalytic and photochemical processes. J. Mol. Catal. 56, pp284-295. Augugliaro, V.; M. J. Lopez-Munoz; L. Palmisano and J. Soria. 1993. Influence of pH on the degradation kinetics of nitrophenol isomers in a heterogeneous photocatalytic system. Appl. Catal. A: General, 101, pp7-13. Kim, D. H. and M. A. Anderson. 1996. Solution factors affecting the photocatalytic and photoelectrocatalytic degradation of formic acid using supported Ti@ thm films. J. Photochem. Photobiol. A: Chem., 94, pp221-229. Pichat, P.; Guillard, C.; Maillard, C.; Amalric, L. and DOliveira, J. C. 1993. In Photocatalytic Purification and Treatment of Water and Air. Edited by D. F. Ollis, H. Al-Ekabi, pp207-223 Amsterdam: Elsevier. Buxton, G. V.; C. L. Greenstock; W. P. Helman and A. B. Ross. 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Ref. Data. 17.~~513-533. Braun, A. M. 1990 In Industrial Photochemistry. Edited by Viriot, M. L.; Andre, J. C.; and Braun, A. M. CPIC-ENSIC: Namcy, France. Vol. A, pp413-422. Abdullah, M.; G. K. C. Low and R.W. Matthews. 1990. Effects of common inorganic anions on rates of photocatalytic oxidation of organic carbon over illuminated titanium dioxide. I. Phys. Chem., 94, pp. 6820-6825. Tseng, J. M. and C. P. Huang. 1991. Removal of chlorophenols from water by photocatalytic oxidation. Water Sci. Technol. 23, pp377-387. Chen, H. Y.; 0. Zahraa and M. Bouchy. 1997. Inhibition of the adsorption and photocatalytic degradation of organic contaminant in an aqueous suspension of Ti02 by inorganic ions. J. Photochem. Photobiol. A: Chem., 108, pp37-44. Haarstrick, A. Kut, M. and Heinzle, E. 1996. Ti02 assisted degradation of environmentally relevant organics compounds in wastewater using a novel fluidized bed photoreactor, Environ. Sci. Technol. 30. p817-824.

pp49-63.

~~2477-2482.

549

D. Chen, M . Sivakumar and A.K. Ray

140.

141

142. 143. 144.

145.

146.

147.

148.

149.

150.

151.

152.

153.

154.

155.

156.

157.

158.

159.

160.

161.

162.

163.

Mills, A. and S. Moms. 1993. Photomineralization of 4-chlorophenol sensitized by titanium dioxide: A study of the annealing the photocatalyst at different temperatures. J. Photochem. Photobiol. A: Chem., 71, pp285-289. Bolton, J. R. and S. R. Cater. 1994. Homogeneous photodegradation of pollutants in contaminated water: An introduction. In Aquatic and Surface Photochemistry. Edited by Hele, G. R. et al., pp467-490. Bolton, J. R.; R. G. Bircher; W. Tumas and C. A. Tolman 1996. J. Adv. Oxid. Technol. 1, pp13. Calvert, J. G. and J. N. Pit& 1966. Photochemistry. Wiley, New York. Serpone, N.; R. Terzian; D. Lawless; P. Kennepohl and G. Sauve. 1993. On the usage of turnover numbers and quantum yields in heterogeneous photocatalysis. J. Photochem. Photobiol. A: Chem.,

Rosenberg, I.; J. R. Brock and A. Heller. 1992. Collection optics of Ti02 photocatalyst on hollow glass microbeads floating on oil slicks. J. Phys. Chem., 96, pp3423-3428. Serpone, N.; G. Sauve; R. Koch; H. Tahiri; P. Rchat; P. Piccinini; E. Peliuetti and H. Hidaka. 1996. Standardization protocol of process efficiencies and activation parameters in heterogeneous photocatalysis: relative photonic efficiencies G. J. Photochem. Photobiol. A: Chem., 94, pp191- 203. Tahiri, H.; N. Serpone and R. L. Mao. 1996. Application of concept of relative photonic efficiencies and surface characterization of a new titania photocatalyst designed for environmental remediation. J. Photochem. Photobiol. A: Chem., 93, pp199-203. Serpone, N. 1997. Relative photonic efficiencies and quantum yields in heterogeneous photocatalysis. J. Photochem. Photobiol. A: Chem., 104, ppl-12. Valladares, J. E. and J. R. Bolton, 1993. A method for the determination of quantum yields in heterogeneous systems: the Ti02 photocatalyzed bleaching of methylene blue. In Photocatalytic purification and treatment of water and air. Ollis, D. F. and Al-Ekabi, H. Eds.. Elsevier: Amsterdam, pplll-120. Bideau, M., Claude], B., Faure, L., and Kazounan, H. 1991 The photooxidation of acetic acid by oxygen in the presence of titaniu dioxide and dissolved copper ions. J. Photochem. Photobiol. A: Chem., 61, pp269-280. Hemnann, J. M.; H. Tahhi; Y. Ait-Ichou; G. Lassaletta; A. R. Gonzalez-Elipe and A. Fernandez. 1997. Characterization and photocatalytic activity in aqueous media of Ti02 and Ag-Ti02 coating on quartz. Applied Catalysis B: Environmental, 13, pp218-229. Modestov, A. D. and 0. Lev 1998. Photocatalytic oxidation of 2,4-dichlorophenoxyacetic acid with titania photocatalyst. Comparison of supported and suspended Ti02. J. Photochem. Photobiol. A: Chem. 112, pp261-270. Minero, C.; E. Peliuetti; S. Malato and J. Blanco. 1996. Large solar plant photocatalytic water decontamination: Effect of operational parameters. Sol. Energy, 56, pp421-428. Minero, C.; E. Peliuetti; S. Malato and J. Blanco. 1996. Large solar plant photocatalytic water decontamination: Degradation of atrazine. Sol. Energy, 56, pp411-419. Ray, Ajay K. 1997. New Photocatalytic Reactors for Destruction of Toxic Water Pollutants, Developments in Chemical Engineering & Mineral Processing, 5, 115-128. Mukherjee, P. S. and Ajay K. Ray, 1999. Major Challenges in the Design of a Large-Scale Photocatalytic Reactor for Water Treatment, Chem. Eng. Technol., 22(3), 253-260. Ray, Ajay K. 1998. Simulation of Photocatalytic Reactor Using CFD", IES Journal Chemical Engineering Issue, 38 (3), 22-27. Penathamby, U. and Ray, Ajay K. 1999. Reactive Flow Modelling on a Distributive Computing Environment, Chem. Eng. TechnoL, 22 (lo), 881-888. Ray, Ajay K. 2000, Design and Development of two large-scale photocatalytic reactors for treatment of toxic organic chemicals in wastewater", in Environmental Reaction Engineering in Pollution Prevention, Elsevier Science, Edited by M. Abraham and R. Hesketh, Chapter 13, ~151-167. Ray, Ajay K. 1998. A New Photocatalytic Reactor for Destruction of Toxic Water Pollutants by Advanced Oxidation Process, Catalysis Today, 44,357-368. Ray, Ajay K. and Beenackers, A. A. C. M., 1998. A Novel Photocatalytic Reactor for Water Purification, AIChE Journal, 44(2) 477-483. Ray, Ajay K. and Beenackers, A. A. C. M.. 1998. Development of a New Photocatalytic Reactor for Purification, Catalysis Today, 40.73-83. Ray, Ajay K. 1999. Design, ModeIIing and Experimentation of a New Large-scale Photocatalytic Reactor for Water Treatment, Chem. Eng. Sci., 54 (15-16), 31 13-3125.

73, ppll-16.

550

heterogeneous photocatalysis in environmental remediation

Documents