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3 1293 0105

 

 

 

 

 

 

i

 

 

 

 

This is to certify that the

thesis entitled

The Photodegradation of 2-Chlorophenol in Aqueous T102
Suspensions: The Effect of Hydrogen Peroxide Addition

presented by

Chin-Chang Wang

has been accepted towards fulfillment
of the requirements for

Master degree in Science

 

Department of Civil and Environmental Engineering

(.2 .

.Mcm H2 Dmfio '

Major professor

 

Date March 24, 1994

 

0-7639 MS U is an Affirmative Action/Equal Opportunity Institution

 

LIBRARY
Michigan State
University

 

 

 

PLACE ll RETURN BOX to romovothb chockout from your rocord.
1'0 AVOID FINES rotum on or botoro doto duo.

DATE DUE DATE DUE DATE DUE

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

MSU Io An Afflrmotlvo Action/Ema) Opportuntty Inotltution
W

THE PHOTODEGRADATION OF 2-CHLOROPHENOL IN AQUEOUS TiOz
SUSPENSIONS: THE EFFECT OF HYDROGEN PEROXIDE ADDITION

By

Chng Wang

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

MASTER OF SCIENCE

Department of Civil and Environmental Engineering

1994

ABSTRACT

THE PHOTODEGRADATION OF 2-CHLOROPHENOL IN AQUEOUS TiOz
SUSPENSIONS: THE EFFECT OF HYDROGEN PEROXIDE ADDITION

By

Chih-Chang Wang

The photocatalytic degradation of 2-chlorophenol (2CP) has been investigated
in aqueous suspensions of TiOz over the wavelength range of 300-390 nm.
The optimum pH range to operate the photocatalytic degradation without
hydrogen peroxide addition is between 5 and 8. With the addition of
hydrogen peroxide the optimum pH range for the photocatalytic degradation
of 2CP is between 3.6 and 5. Bubbling air or pure oxygen into UV/Ti02 system
does not increase the rate of degradation of 2CP to a great extent. The
degradation rate of 2CP was faster when both hydrogen peroxide and oxygen
were present. At 100% of light intensity (3 mW/cmz) the effect of adding
hydrogen peroxide on photocatalytic degradation rate was significant.
However, at the 35% of light intensity the addition of hydrogen peroxide did

not significantly increase the photocatalytic degradation rate.

DEDICATION

To My Parents

iii

ACKNOWLEDGMENTS

I would like to express my sincere appreciation to Dr. Simon H. Davies for
his valuable guidance and encouragement throughout the course of this
study.

I also would like to thank Dr. Susan J. Masten and Dr. Craig S. Criddle for
their helpful suggestions.

The financial support from Great Lakes and Mid-Atlantic Hazardous
Substance Research Center is gratefully acknowledged.

In addition, I would like to thank Dr. Paul R. Loconto and Yan-Lyang Pan
for sharing their technical expertise in HPLC.

Appreciation is also extended to my parents who provided some financial
support.

Special thanks go to my wife, Lily, for her patience, understanding, steady

encouragement, and many sacrifices throughout this study.

iv

TABLE OF CONTENTS

List of figures ............................................................................................................... viii
List of tables ...................................................................................................................... x
I. Intruduction ............................................................................................................ 1
1.1 The reason for studying 2-chlorophenol .................................................. 1
1.2 The process to decompose 2-chlorophenol .............................................. 4
1.3 Outline of this study ...................................................................................... 5
11. Background ............................................................................................................. 7
2-1 Preperties of TiOz .......................................................................................... 7
2.2 Photoprocesses on TiOz surface .................................................................. 8
2.3 Photodegradation of chloroaromatic and chloroalphatic
compounds in TiOz suspensions ............................................................. 12
2.4 pH effect ......................................................................................................... 15
2.5 light intensity effect ..................................................................................... 16
2.6 ()2 effect .......................................................................................................... 17
2.7 Fenton’s Reagent and Fenton-like Chemistry ....................................... 22
2.8 Hydrogen peroxide with UV radiation over Ti02 suspensions ........ 23
2.9 Kinetics studies ............................................................................................. 25
III. Materials and methods ........................................................................................ 28
3.1 Experimental equipment ........................................................................... 28
3.1.1 Photochemical reactor .................................................................... 28
3.1.2 Ultraviolet light source .................................................................. 30
3.1.3 Experimental scheme ...................................................................... 30
3.1.4 pH control system ............................................................................ 33
3.1.5 High Performance Liquid Chromatograph ................................ 33
3.1.6 Spectrophotometer .......................................................................... 34
3.1.7 Sample collection ............................................................................. 34
3.1.8 Nonlinear fit computer program ................................................. 34
3.2 Experiments .................................................................................................. 35
3.2.1 Materials ............................................................................................ 35
3.2.2 Chemical preparations .................................................................... 35
3.2.3 Standard calibration curve ............................................................. 36
3.2.4 pH calibration ................................................................................... 36
3.2.5 Batch experiments-UV/ Titanium dioxide system ................... 36
3.2.6 Batch experiments-UV/ Titanium dioxide/ Hydrogen peroxide
system ................................................................................................. 37

3.2.7 Effect of pH ........................................................................................ 37

3.2.8 Effect of light intensity .................................................................... 38

3.2.9 Potassium ferrioxalate actinometry ............................................. 38
3.2.10 Effect of illuminated method ........................................................ 38
3.2.11 Effect of oxygen concentration ...................................................... 39
3.2.12 Effect of Fe2+ ..................................................................................... 39
3.2.13 Buffer system .................................................................................... 39
IV. Results and discussion ......................................................................................... 40
4.1 UV radiation combined with air, H202. or Ti02 .................................. 42
4.1.1 UV/ Air ............................................................................................... 42
4.1.2 UV/ H202 ........................................................................................... 43
4.1.3 UV/TiOZ ............................................................................................ 44
4.1.4 Comparison of UV radiation combined with air, H202, or
Ti02 ..................................................................................................... 46
4.2 UV/Ti02 combined with air or pure oxygen ........................................ 46
4.2.1 UV/ Ti02/ Air .................................................................................... 46
4.2.2 UV/Ti02/02 ..................................................................................... 47
4.2.3 Effect of 02 concentration on UV/Ti02, UV/ Ti02/ Air, and
UV/ Ti02/02 system ....................................................................... 48
4.3 UV / Ti02 combined with H202 ................................................................ 52
4.3.1 UV/Ti02/ H202 ................................................................................ 52
4.3.2 Comparison of UV/ H202 and UV/ Ti02 / H202 system ......... 52
4.4 Photocatalytic systems under intermittent illumination method ...53
4.4.1 UV / Ti02/ 02 system under intermittent method ................... 53
4.4.2 UV / Ti02/ H202 under intermittent method ............................ 54
4.4.3 Effect of intermittent method on UV/Ti02/02 and
UV/ Ti02/ H202 system .................................................................. 54
4.5 UV/Ti02 under intermittent method and pH control ....................... 55
4.6 UV / Ti02 under buffer system .................................................................. 56
4.7 Effect of pH on photocatalytic systems .................................................... 58
4.7.1 UV/ Ti02/ Air / NaHC03 with and without H202 under pH
control ................................................................................................ 58
4.7.2 Effect of pH on UV/ Ti02 / Air/ N aHC03 with and without
H202 ................................................................................................... 60
4.8 Effect of Fe2+ on photocatalytic systems ................................................. 62
4.8.1 UV/Ti02/Air/NaHC03/Fe2+ with and without H202 ,,,,,,,,, 62
4.8.2 Effect of adding ferrous chloride ................................................... 62

vi

4.9 Effect of light intensity on photocatalytic systems ................................ 64
4.9.1 UV/Ti02/ Air/NaHC03 with and without H202 under

different light intensity ................................................................... 64

4.9.2 Effect of light intensity .................................................................... 65

4.10 Summary of reaction rate constants ........................................................ 68

V. Conclusions and Recommendations ............................................................... 69
5.1 Conclusions ................................................................................................... 69

5.2 Recommendations ...................................................................................... 70

List of references ............................................................................................................. 71
Appendix A. Experimental data ............................................................................... 77

vii

LIST OF FIGURES

Figure Page
2.1 Electron transfer processes at the semiconductor-water interface ............... 8
2.2 Model for reaction between 02, diffusing from solution to particle

2.3

2.4

3.1
3.2
3.3

4.1

4.2

4.3

4.4

4.5

4.6

4.7

surface, and highly mobile electrons. The reactions take place when the
02 molecule comes within an electron-transfer distance 8 from the
surface ...................................................................................................................... 19

Model of Figure 2.2 modified by a limitation of the reaction with 02
molecules to electrons at a distance 8 from the interface ............................. 20

Model for reaction in which the 02 molecules have to diffuse from
solution to surface sites, where a trapped electron resides. Only oxygen

molecules within a distance 6 can react with trapped electrons ............... 21
Schematic of Supermix® Photochemical Reactor ........................................ 29
Structure of UV lamp .......................................................................................... 31
Experimental Scheme .......................................................................................... 32

Effect of 02 concentration on photodegradation of 2-chlorophenol in
UV/Ti02, UV/ Ti02/ Air, and UV/ Ti02 / 02 system ..................................... 49

Effect of oxygen concentration on the photocatalytic degradation rate of 2-
chlorophenol ......................................................................................................... 51

Comparison of different illumination methods in same oxidation
system ...................................................................................................................... 55

The effect of HC03' on the photocatalytic degradation rate of 2-
chlorophenol ......................................................................................................... 58

Effect of initial pH on UV/ Ti02/ Air / NaHC03 system with and without
H202 ........................................................................................................................ 60

Effect of Fe2+ on UV/Ti02/Air/NaHC03 with and without H202 ,,,,,,,,, 63

Effect of light intensity on UV / Ti02/ Air/ NaHC03 system with and
without H202 ........................................................................................................ 65

viii

4.8 Effect of light intensity on the rate of photocatalytic degradation of 2-
chlorophenol with and without adding hydrogen peroxide addition

(where 100%=3 mW/cmz) .................................................................................. 67

ix

LIST OF TABLES

Table Page
1.1 Physical properties of the monochlorophenols ............................................... 2
1.2 Production and use of the chlorophenols ......................................................... 3
2.1 Photocatalystic Reaction Scheme ...................................................................... 10
2.2 Half-lives for various Chloroaromatics ........................................................... 13
4.1 The initial conditions of experiments .............................................................. 41
4.2 The results of 2-chlorophenol degradation in UV/Ti02 system under

4.3

4.4

4.5

intermittent method at different pH during 26 min illumination ........... 56

The results of 2-chlorophenol degradation in UV/Ti02/ Air/ NaHC03

with and without H202 at different initial pH .............................................. 59
The results of 2—chlorophenol degradation in UV/Ti02/ Air/ NaHC03

with and without H202 at different light intensity ..................................... 64
Initial quantum yield of 2-chlorophenol ......................................................... 68

Chapter 1

Introduction

1.1 The reason for studying 2-chlorophenol

The widespread utilization of chlorinated aromatic compounds as pesticides
and herbicides is attracting increased concern (1). The occurrence of highly
chlorinated organic compounds, such as polychlorinated biphenyls and
chlorobenzene is wide spread in the environment (2). These compounds are
important from a health perspective due to their carcinogenicity or potential
carcinogenicity (3,4) and their persistence in the environment (5). This has led
to a better understanding of the influences of chemicals that contaminate
water which endangers both the environment and human health.
Chlorinated organics are abundant pollutants and some of them withstand

biodegradation. Among them, monochlorophenols are moderately toxic to

aquatic life. Chlorophenols induce disagreeable taste and odour in water and

accordingly impair the flavor of fish even at ppb (ug-kg’l) levels. Several of

the important physical properties of monochlorophenols are summarized in

Table 1.1 (6).

 

 

Table 1.1 Ph ical ro r 'es f m nochlor henols (6)
Parameter 2-chlorophenol 3-chlorophenol 4-chlorophenol
Chemical OH (:1 OH
structure ‘0 a
Cl
Molecular mass 128.6 128.6 128.6
Physical state liquid solid solid
Color colorless white white
Melting point (°C) 9.3 33 43
Boiling point (°C) 175 214 217
Density d? 1.257 di‘ 1.245 d? 1.383
Vapor pressure 1 mm Hg at 1 mm Hg at 1 mm Hg at
12.1 °C 44.2 °C 49.8 °C
Water solubility 28.5 26 27.1
(at 20 °C, g/ 1)
pKa 8.48 9.02 9.38

 

There are two important techniques to synthesize chlorophenols: direct

chlorination of phenol and alkaline hydrolysis of polychlorobenzenes (7).

When gases chlorine is passed into molten phenol at temperatures between

50 and 150 °C, both 2-chlorophenol and 4-chlorophenol are formed. The

method of production and principal uses of the chlorophenols are presented

 

 

in Table 1.2.
Table 1.2 Production ad use of the chlorophenols (6)
Production Use
2-CP Usually synthesized through Intermediate in the synthesis of

chlorination of phenol: may be
synthesized from 1,2-
dichlorobenzene by alkaline
hydrolysis, or from diazotized _o_-
chloroaniline.

May be prepared by hydrolysis of
1,3-dichlorobenzene, from r_n_-
chloroaniline through the
diazonium salt, or by oxidation of
p-chlorobenzenic acid with
copper(H) oxides as the oxidant

0n large scale by chlorination of
phenol; may be synthesized from
p-chloroaniline, from p-
nitrosophenol, by selective
reduction of chlorobromophenols
or by alkaline hydrolysis of 1,4-
dichlorobenzeze.

higher chlorophenols and
phenolic resin

Chemical intermediate

Intermediate in production of
higher chlorinated phenols, and
in synthesis of dyes and drugs
(denaturant for alcohol;
selective solvent in refining of
mineral oils.

 

Industrial waste discharge is the principal source of water pollution by

chlorophenols. These compounds are produced as end- or by-products in

chemical synthesis or may be used as starting materials as well. Consequently,

that is the reason to conduct these experiments to study chlorophenols and
find out efficient methods to reduce the amount of them in environment. 2-
chlorophenol was chosen in this study because it is a liquid which is easier to

be prepared as a reactive solution.
1.2 The process to decompose 2-chlorophenol

There are several treatment alternatives to decompose chlorophenols in
wastewater. Among them, biological treatment processes are widely used
methods because of their effectiveness and low cost. Biotransformation by
microorganisms in water and soil is considered to be the major mechanism
by which chlorophenols are degraded in the environment (6). It has been
shown that the rate of biodegradation is largely dependent on the microbial
populations present, and the characteristics of the environment, like pH,
oxygen level, nutrient status and temperature. The most noticeable drawback
of the biodegradation is that when concentration of the pollutants in
wastewater is too high, the microorganism can not function properly. In
many cases, the concentration of chlorophenols in wastewater is high
enough to be toxic to microorganisms. Another reason for the unsuitability of
the biological treatment process are that it always needs long detention time
or long start-up time to acclimate the microorganisms to the waste. In the
conventional wastewater treatment processes, activated carbon adsorption
which is a non-destructive process is used to remove contaminants from
wastewaters, but it only concentrates the contaminants rather than destroying
them. Thermal destruction could be used to treat aqueous wastes. However,
the cost for treatment of diluted wastes is generally high. Alternative

technologies which can efficiently decompose the chlorophenols in

wastewaters are needed. Advanced oxidation processes (AOP) are a possible

technologies to solve this problem.

AOPs are defined as those technologies that involve the generation of highly
reactive radical intermediates, particularly the hydroxyl radical (OH), at

ambient temperature (8). The typical AOP systems utilize combination of

strong oxidants, such as hydrogen peroxide and ozone, to conduct the process.

The combinations used include 03/ UV, 03/ H202, 03/ H202/ UV, Ti02 / UV,
Ti02 /H202 / UV, Ti02/ UV / Fe2+, and Ti02/ UV/ H202/ Fe2+. Ti02 combined
with UV light and hydrogen peroxide was chosen as major experimental

system in this study.

1.3 Outline of this study

The objective of this study is to improve the understanding of the
photodegradation of chloroaromatic compounds in titanium dioxide

suspensions, and also determine the feasibility of using Ti02/ UV,
Ti02/UV/H202, Ti02/UV /1=e2+, or Ti02/UV/H202/Fe2+ to degrade 2-

chlorophenol in a aqueous suspensions. The specific objectives are:

1. to determine an optimum condition to operate the AOP system in a

effective way.

2. to study the effect of experimental parameters such as pH, concentration

of H202, and light intensity on the photocatalytic degradation of 2-

chlorophenol.

u

and

3. to study the influence of the presence of air and oxygen on the Ti02

induced photodegradation of the target compound.

Chapter 2

Background

2.1 Properties of Ti02

Ti02 is an n-type semiconductor and has a band gap energy of 3.2 eV (9). At
pH 7, the conduction band is located at 2.6 V vs. NHE (10), so it has the
potential to oxidize a great number of organic compounds (10). Ti02 has been
favored as a photocatalyst for treating aqueous wastes because of the strong
oxidizing potential of its valance band holes and because it is not
photodecomposed. In an n-type semiconductor, the reactive species is the
hole, cogenerated with an electron upon absorption of a photon and carrying
the major part of the light quantum energy. This excess free energy is the
driving force for the oxidation of the organic and inorganic molecules that
react at the surface of such a particle. Other potential photocatalysts, for
example, ZnO will dissolve during photochemical reactions and generate

undesirable products, such as Zn2+. Two further advantages of using Ti02 as

U

photocatalyst are (i) it is a true catalyst because it is not consumed during the
experiment and (ii) it can be easily immobilized, so the treatment process can

be adapted for continuous operation (11).
2.2 Photoprocesses on Ti02 surface

The use of T102 as a photocatalyst for the degradation of organic pollutants
shows considerable promise. The photo-induced processes in semi-conductor
suspensions have been extensively reviewed (9, 12). When a semiconductor
is illuminated with light of energy greater than its bandgap energy, an
electron is transferred to the conduction band leaving a vacancy (”hole”) in
the valance band. The electron and hole can migrate to the surface of the
semiconductor. At the interface, electron transfer can occur either from the
conduction band to an acceptor, A, in solution or from a donor, D, in solution

to the valence band. These processes are illustrated in Figure 2.1 (13).

A
h A
‘ D - Electron donor
Ti02 h‘] A - Electron acceptor

CB - Conduction band
VB 6-)
v D+

VB - Valance band
D

 

 

Figure 2.1 Electron transfer processes at the semiconductor-water interface.

Turchi and Ollis (1990) proposed that the hydroxyl radical, -OH, is the primary
oxidant in the photocatalytic system (14). The photocatalytic reaction scheme

for the system proposed by Turchi and Ollis is given in Table 2.1.

When a semiconductor is irradiated with light of energy greater than its
bandgap energy, an electron is excited into conduction band leaving a vacancy
(”hole”) in the valence band (reaction 1). The electron and hole can migrate to
the surface of the semiconductor. If reactive species (traps) are not present at
the surface of the semiconductor then the electron-hole pair will recombine
(reaction 5) and the energy absorbed will be dissipated as heat. The efficient
trapping of the hole and electron at the surface (reaction 6-8, 13, and 14) is a
key factor in achieving efficient photochemistry. It is important that the
electron acceptor and donor be adsorbed on the Ti02 surface, since the initial
reactions occur in the picosecond time domain (9). In water the solvent, itself,
is a potential electron acceptor. However, since water is not easy to reduce,
where oxygen is present, oxygen is the most likely electron acceptor (reaction
8a and 8b). Adsorbed oxygen is reduced to form superoxide, 0‘2. Superoxide is
unstable and can be reduced to form hydrogen peroxide (reaction 13) and the
perhydroxyl ion (reaction 14). On the Ti02 surface, it has been suggested that
the holes in the valence band appear to react with water to produce not only
the hydroxyl radicals, but, possibly, another unidentified weaker oxidant (15).
Holes are though to react with surface OH‘ and H20 groups to form -OH
(reaction 6). An adsorbed organic molecule, R1, ads may be directly oxidized
by a hole (reaction 7); however, this reaction is not believed to be significant,
because of the lack of reactivity that is observed in organic solution (14). The

oxidation of organic compounds is though to occur

10

T 12.1Pht a1 'cR 'n heme

 

Excitation
Ti02 -—“"——)e’ + h+ (1)
Adsorption
of: + Ti" + H20 4: OLH‘ + TiIv - OH‘ (2a)
1i” + H20 4: Ti" - H20 (2b)
site + R1 in Rm, (3)
0H . +11” t: Ti‘Vu-OH- (4)
Recombination
c“ + h” —-> heat (5)
Trapping
Ti” —0H' + h” 4: Ti‘V---0H- (6a)
TiIv - H20 + h+ {z} Ti‘v- . .OH . +H“ (6b)
Rm, + h+ 4:» RI“, (7)
Ti" + e‘ 4:} Tim (8a)
Tim +02 4: Ti” — 0; (8b)
Hydroxyl Attack
CaseI Tiw---0H-+R1,.a. —> TiIv + Rm (9)
Case H OH ' +Rt,.a. ‘9 R2,“). (10)
CaseIII 'ri“’---0H-+Rl —>Ti“’ +122 (11)
Case IV OH-+R1 -9 R2 (12)
Reaction of other radicals
c“ + '11” - o; + 2(H+) 4: TiN(H202) (13)
Ti” — .o; + (W) 4: Ti”(02H) (14)

(H202) +0H- 4:) (.0211) + (H20) (15)

 

11

as a result of the reaction of the molecule either with the adsorbed ~0H
(reactions 9 and 11) or with ~0H that is not bonded to the surface (reactions 10
and 12). Another reaction that may occur in the system is the scavenging of

-0H by hydrogen peroxide (reaction 15).

It is proposed that the hydroxyl radical ,-0H, is the primary oxidant during the
photocatalyzed degradation experiments in irradiated Ti02 suspensions (14).

The formation of -OH radicals can be achieved by two routes: (a) via reaction

of valance band holes with either adsorbed H20 or with surface OH’ groups

(reactions 1, 16, 17)

1133 + H20(ads) -> OH + H+ (16)
hi,la + OH' (surf) —9 -OH (17)

or (b) via the H202 reaction from 0', (16).

It is generally accepted that adsorption of oxygen inhibits the electron/ hole

recombination process by trapping the conduction band electron as a

superoxide ion, 0;.

02 + cgm —> 0', (18)

Subsequent to reaction 18 , H202 can be formed from 0;, via reactions 19-22.

0; + 11* —> 1102 - (19)
Ho2 -+H02- —+ 11,02 +02 (20)

12

o; + no, —) 110; + 02 (21)
H0; + H” -) I1202 (22)

Cleavage of H202 by any of reactions 21-23 yields -OH radicals.

H202 + e33 --) OH + OH' (23)
H202 + O; —-) -0H + OH' + 02 (24)
H202—”4+2 - 0H (25)

In order for the photooxidation process to occur it is necessary to avoid
accumulation of the electrons or holes in the particles. Because the
accumulation would increase the recombination rate and lower the quantum
yield. Thus, the electrons and holes must be removed promptly by an electron
acceptor or donor at the surface. The faster this removal, the lower the
concentration of electrons an holes in the particles. The lower the steady-state
concentration of electrons on the particles, the lower the recombination losses

(16).

2.3 Photodegradation of chloroaromatic and chloroalphatic

compounds in Ti02 suspensions

The photodegradation of chloroaromatic compounds in Ti02 suspensions
has been extensively reviewed (17). All the compounds listed in Table 2.2

could be completely mineralized.

The half-lives for the disappearance for these compounds ranged from 10 to

90 minutes.

13

Table 2.2 Half-lives for various Chlgrgargmatics

 

Compound Concentration (ppm) t1 / 2 (min)
4-chlorophenol 6 14
3,4-dichlorophenol 18 45
2,4,5-trichlorophenol 20 55
Pentachlorophenol 12 20
Chlorobenzene 45 90
1,2,4-trichlorobenzene 10 24
3,3-dichlorobiphenyl 1 10

 

catalyst concentration 2 g/ L; aerated solution; pH 3 except for chlorobenzene;

pH 2.5 for chlorobenzene; suspension irradiated with light>330 nm (17).

The photocatalyzed mineralization of many chlorinated aliphatic and
aromatic hydrocarbons, as well as other organic contaminants, in drinking
water has been successfully demonstrated (16). The total mineralization of
many of these compounds to C02, Cl’ and H20 via heterogeneous
photocatalysis mediated by illuminated Ti02 has been demonstrated. Direct
photolysis of these compounds by UV light generally only leads to partial
degradation (16).

Al-Sayyed, D’Oliveira, and Pichat (18), have demonstrated that in the absence
of a semiconductor, 4-chlorophenol is destroyed completely in 100 min (when
irradiated with light of wavelength greater than 290 nm, i.e. within the
absorption band of 4-chlorophenol). In the presence of Ti02, a rapid

degradation occurs and the concentration of pollutant is below detection

limits within 60 min. The results show that the photodegradation in a Ti02

14

suspension is more efficient than the direct photolysis. Al-Sayyed et al. also
found that the rate of a photocatalytic reaction increased linearly with the

amount of catalyst up to a level corresponding to complete adsorption of the

incident light by Ti02.

Al-Ekabi et al. in 1991 (19), showed that the degradation of both 2,4-
dichlorophenol and pentachlorophenol was very efficient ,and in 12 min it

was possible to degrade 95% of 2,4-dichlorophenol in an aerated aqueous Ti02

suspension.

Photolysis in the presence of Ti02 can detoxify PCBs. Oliver and Carey in 1985
(15), irradiated a 30 ppb Aroclor 1254 solution containing 0.5% anatase with

350 nm light for 1 hour. They found that the growth of the green algae
W was virtually the same in the irradiated Aroclor

solution and the distilled water control. Algae growth in the unirradiated

Aroclor solution was about 100 times slower than in the control.

Pruden and Ollis (20), in 1983, showed that the chloroform is degraded readily

in illuminated suspensions of Ti02 according to the following stoichiometry:

2CHC13 + 21120 + 02 J» 2C02 + 6HC1 (26)

Their results showed that in the presence of both Ti02 and near-UV light,
CHC13 was rapidly dehalogenated, while in the presence of either UV light or
Ti02 alone, no CHC13 disappearance was observed.

15

2.4 pH effect

A key factor that affects the degradation rate of chemicals in Ti02 suspensions
is the pH. The results of Al-Ekabi et a1. (16) show that the rate of degradation
of 4—chlorophenol at pH 2.5 is 4 times less than that at pH 5.8. Hydrochloric
acid appears to inhibit the degradation of 4-chlorophenol. As the addition of
chloride (as KCl) does not inhibit the reaction, it is suggested that the
protonation of the Ti02 surface is responsible for the inhibition seen at low
pH (16). Ollis et a1. (21) has also suggested that at low pH the adsorption of H4”
may be responsible for the inhibition of Ti02 mediated photodegradation of
chlorinated hydrocarbons. The inhibition seen at low pH may be due to
changes in adsorption/ desorption characteristics of the reactants and/ or other
species with pH. A shift in surface potential of the semiconductor as result of

H+ ion sorption may also important.

Kormann et a1. (22) showed that simple anions such as Cl' and HC03'
decrease the photoefficiency of the catalyst at low pH where they can be
absorbed to the positively charged T102 particle surface (szpc=6.25 for P25
Ti02). However, the same anions have a negligible affect on the measured
photoefficiency at high pH since their adsorption to reactive surface sites is
prevented by electrostatic repulsion from the negatively charged Ti02

particles.

D’Oliveira et al. (23) have studied the effect of pH variation on the Ti02
induced photodegradation of 3—chlorophenol. At pH 10.8, 4.5 and 2.5 the time
for 99% disappearance of 3-chlorophenol was 90, 115 and 190 minutes,

respectively. The increase in the rate of photodegradation with increasing pH

16

above can be attribute to the increased number of OH' ions at the surface of
Ti02, since -OH radicals can be formed by trapping photoproduced holes.

Similarly, the decrease at low pH can be explained by the lack of OH' ions.

Al-Sayyed et al. in 1991 (18) demonstrated that the rate of photocatalytic
degradation of 4-chlorophenol solutions was independent of pH in the range
3.4-6.0. Similar observations were also reported by Al-Ekabi et al. in 1989 (16).

2.5 light intensity effect

Kormann et al. (22) found that the rate of chloroform degradation is a
nonlinear function of light intensity. And the quantum yield of the reaction
increases with decreasing light intensity. It was found that the rate of

chloroform degradation was proportional to half power of light intensity, i .e.,

_d[C:l-ItCl§] = kIl/Z (27)

D’Oliveira et al. (23) found that the rate of photodegradation of 3-
chlorophenol was proportional to the radiant flux (or light intensity) for
values smaller than ca. 20 chm'z. Above this value, it is proportional to its
square root, which means that the recombination of the photoproduced
charges predominates and accordingly, the initial quantum yield of the

degradation process decreases with increasing light intensity.

Al-Sayyed et al. (18) found that the relationship between the
photodegradation rate and light intensity of 4-chlorophenol was nearly linear

at low light intensity, but it was evident that a leveling off occur at higher

17

light intensity. The rate of 4-chlorophenol disappearance became proportional
to the square root of the light intensity. The results suggest that the
recombination of photoproduced charges is predominant over oxidation of

organic molecules at high light intensities.

2.6 02 effect

In many studies on photocatalysis at semiconductor surfaces it has been
found that in the absence of oxygen the reaction does not occur (15, 17). The
role oxygen plays is not fully understood. It is believed that adsorbed oxygen
acts as an electron acceptor (17), as such the presence of oxygen tends to retard
hole-electron recombination. The presence of oxygen may also enhance the
formation of hydrogen peroxide (see reaction 19-22) and eventually ~OH (see
reaction 23-25).

D’Oliveira et al. (23) found that oxygen (or another electrophilic species) was
required for photocatalytic oxidations of 2- and 3-chlorophenol. From
measurements carried out for the system 02-UV-illuminated Ti02 the
existence of several types of negatively charged adsorbed oxygen species has
been postulated: 02‘, 03", 0', and 022' (23). The formation of H02- radicals
from 02' and H+ giving rise to H202 has been proposed as another way of
producing t0H radicals. This way would be favored at low pH.

Al-Ekabi et al. (19) showed that the introduction of oxygen led to an 80%
increase in the degradation rate of 2,4-dichlorophenol compared to the air
saturated solution. Under nitrogen the reaction was 75% slower than in air.

The data clearly indicate that oxygen plays a crucial role in the photocatalytic

 

18

degradation of 2,4-dichlorophenol. They mentioned that introducing small
bubbles of either air or oxygen may improve the mixing process inside the
reactor and thus increase the degradation rate. However, it is not yet clear
whether the role of oxygen is simply to accept the conduction band electron
and thus prolong the lifetime of the photogenerated ”hole” or whether it

directly participates in the overall degradation process.

Gerischer et al. (24) studied the role of oxygen in the photooxidation of
organic molecules on semiconductor particles, and proposed two models to
explain the kinetics of the reaction between dioxygen and an electron on the

semiconductor particle.

They assumed that oxygen is the only electron acceptor available initially. The

primary reaction of electrons reaching the surface is:

6;“, + 02 -—> 0; (28)

Because the solubility of oxygen is relatively low and O2 is not strongly
adsorbed on semiconductor surfaces in contact with aqueous electrolyte, a

high proportion of photoproduced electrons recombine with holes.

In case A, an electron moves freely in or on the particle and reacts anywhere
on the particle surface with oxygen. The assumption, made for this process, is
that the reduction can occur anywhere on the surface as long as an 02
molecule is within the electron transfer distance 6. This requires that the

electron mobility be high. This model is shown in Figure 2.2.

19

 

Figure 2.2 Model for reaction between 02, diffusing from solution to particle
surface, and highly mobile electrons. The reactions take place when the 02

molecule comes within an electron-transfer distance 6 from the surface.

Several assumptions in this model are unrealistic. One is the high mobility of
the electrons, which guarantees that all 02 molecules find a reaction partner
during an encounter with the particle. This should be corrected by
introducing a fraction of electrons on the particle that can interact. It is also
postulated that electrons in a spherical shell underneath the particle surface
within an interaction distance a from the interface take part in the reaction.

This is shown in Figure 2.3.

 

20

 

Figure 2.3 Model of Figure 2.2 modified by a limitation of the reaction with 02

molecules to electrons at a distance a from the interface.

A further correction takes into account the trapping of electrons in the bulk,
which immobilizes part of them and leads to another decrease of the apparent

rate constant of electron transfer.

In case B, an electron is trapped at or very close to the surface and reacts with

oxygen from its trapped state (see Figure 2.4). The modified reaction is:

6;.” + 02 —> 0; (29)

 

21

 

Q 02

Figure 2.4 Model for reaction in which the 02 molecules have to diffuse from

solution to surface sites, where a trapped electron resides. Only oxygen

molecules within a distance 6 can react with trapped electrons.

After the analysis of case A and case B, their results showed that case B

appears to be the most realistic.

Gerischer and Heller (24) suggest that the rate of electron transfer to oxygen
(reaction 8) is the decisive factor in determining the overall efficiency of the
photocatalytic process. They also suggest one means of improving the

efficiency of the process is to catalyze the electron transfer to oxygen.

22

2.7 Fenton’s Reagent and Fenton-like Chemistry

In 1876, H.].H. Fenton published the first account of the reaction between

ferrous ion and hydrogen peroxide.

Metelitsa (25) proposed the following mechanism for the reactions occurring

in the Fenton’s system.

Ft:2+ + H202 —) Ft:3+ + OH' + H0- (30)
Fe“ + Ho —9 Fe3+ + 0H” (31)
11202 + H0 —9 HO2 ~+H20 (32)
Fe2+ + H02- —-) Fe“ + H0; 4: H202 (33)
Fe” +H02- 4:) Fe2+ +H+ +02 (34)
Fe3+ +11202 —) Fe2+ +Ho2 ~+H+ (35)

Metelitsa conclude the following refers to the pure system. In acid solution
with an excess of hydrogen peroxide, reactions 34 and 35 can be neglected and
reactions 30, 32, and 33 predominate. With an excess of ferrous ions, reactions
30 and 31 predominate. When the concentration of ferrous ions are
approximately equal to the concentration of hydrogen peroxide, only

reactions 30, 31, 32, and 33 are important.

Mertz and Waters (26) demonstrated that the oxidation of organic compounds
by Fenton’s reagent could be occurred by both chain and non-chain reaction
mechanisms. The rate limiting step is the formation of hydroxyl radicals. The

mechanism is the following.

23

Ft:2+ + H202 -) Fe3+ + OH' + H0- ; chain initiation (36)
R — H + H0- —) R - +1120 ; chain propagation (37)
R - +11202 —) ROH + H0- ; chain propagation (38)
Fe“ + H0 —) Ft:3+ + OH' ; chain termination (39)
R - +I-IO- —) ROH ; non-chain termination (40)
2R- —> Products ; non-chain termination (41)

The classical Fenton's reagent chemistry is conducted in acid solution with
relatively high hydrogen peroxide to ferrous ion ratios. The purpose of this is
to prevent the precipitation of the iron as iron hydroxide. In the neutral pH
range, researchers have to give considerations to the autooxidation and the

type of buffer utilized.

2.8 Hydrogen peroxide with UV radiation over Ti02 suspensions

Based on the reaction scheme presented in section 2.2, the addition of
hydrogen peroxide would generate -0H radicals (see reaction 19-22). As it is
widely believed that -0H radical is responsible for the oxidation of organic
compounds in illuminated Ti02 suspensions, one would expected that the
addition of hydrogen peroxide would accelerate the photodegradation of
soer compounds in these system. Also, the addition of hydrogen peroxide
would be expected to enhance the rate of oxidation of organic compounds, as
its decomposition generates oxygen which also enhances the rate of

degradation of organic compounds (27).

Tanaka et al. (28) studied the degradation of trichloroethylene as a function of

illumination time, and the effect of added hydrogen peroxide on the

24

degradation rate. Their results showed that addition of 4 mM hydrogen
peroxide accelerated the degradation rate by 6 to 8 times. In the absence of
Ti02, only 5% was degraded after illumination for 10 minutes. When
trichloroethylene alone was illuminated virtually, no degradation was

observed within 1 hour.

They concluded:

0 The optimal concentration of H202 was 4x10‘3 to 1.2x10'2 for 3.5x10’4 M of

trichloroethylene.

0 The photodegradation rate in the absence of Ti02 increased with the

concentration of H202.

0 The same effects of added H202 were observed when W03 was employed
as a catalyst instead of Ti02.

Harada et al. (29) have shown that the addition of hydrogen peroxide
enhanced the photocatalytic degradation of dimethyl-2,2-dichlorovynyl
phosphate (DDVP) and dimethyl-2,2,2-trichloro-1-hydroxyethyl phosphate
(DEP). The effect of the hydrogen peroxide depends on its concentration.
Control experiments conducted in the absence of Ti02 or H202 showed that
this effect was not merely the sum of the effect of illuminated H202 and that
of illuminated catalyst. They found the addition of 1.2x10‘2 M H202 increased
the rate of degradation of DDVP by 10 fold.

25

Auguliaro et al. (30) studied the influence of hydrogen peroxide addition on
the photodegradation of phenol in Ti02 suspensions. They found that the

addition of hydrogen peroxide enhanced the rate of degradation significantly.
The photodegradation rate increased at H202 concentrations up to 5x10’3 M,
above that concentration the rate of reaction was independent of H202
concentration. The photodegradation rate was greatest when both oxygen and

hydrogen peroxide were present.

Pacheco and Holmes (31) studied the influence of H202 addition on the

photocatalytic oxidation of salicylic acid in the presence of Ti02. They found
the addition of 100 ppm (2.9x10'3 M) H202 increased the initial degradation
rate by 8 fold. When 300 ppm (8,8x10‘3 M) H202 was added the salicylic acid

degradation rate was 10 times greater than with no added peroxide.
2.9 Kinetics studies

The rate of the photocatalytic degeneration of chlorinated compounds
depends upon various parameters including initial concentration of the
compound, pH, light intensity, temperature, wavelength, mass and type of

photocatalyst, and type of photoreactor.

Kormann et al. (22) showed that the degradation (oxidation) of a general

chlorinated hydrocarbon, CxHyClz, can be described by the following

stoichiometry.

 

 

C,H,c1, +(x+ y;z)02 —) rtco2 +in+ + 20‘ +(Y;Z)H,o (42)

 

26

The C02 will be present as HC03' (pK31=6.3) and C032' (pKa2=10.3). The
heterogeneous rate of photodegradation of chlorinated compound can be

described in terms of a second order kinetic expression.

_d[c,H,c1,]

dt = k[0,],,,[cxii,c1, 1“,, (43)

1102],”: the concentration of oxygen adsorbed on the Ti02 surface
[C,H,Clz]”,: the concentration of chlorinated compound adsorbed on the

Ti02 surface

Harada et al. (29) found that the photodegradation reaction of DDVP and DEP

follows first-order kinetics in illuminated Ti02 suspensions. Similarly, Al-
Ekabi et al. (19) found that the photocatalytic degradation of 2,4-
dichlorophenol can be described by first order kinetics (up to 80%
degradation).

Pruden and Ollis (20) studied the degradation of chloroform under UV-Ti02
suspensions. They showed that the reaction rates for both chloroform

degradation and chloride production followed first-order kinetics.

Al-Ekabi et al. (16) studied the photocatalytic degradation of phenol, 4-
chlorophenol, 2,4-dichlorophenol, and 2,4,5-trichlorophenol over Ti02. The
results indicate that the degradation of the four phenols in aerated aqueous
solutions was a function of irradiation time, and that the reaction kinetics
were approximately first-order to a high degree of degradation (>85% for

phenol).

27

Similarly, Okamoto et al. (32) in the study of heterogeneous photocatalytic
decomposition of phenol over Ti02 powder showed that the degradation

reaction followed the first-order kinetics.

Chapter 3

Materials and Methods

3.1 Experimental equipment

3.1.1 Photochemical reactor

The photochemical reactor used to conduct the studies described in this work
is illustrated in Figure 3.1. It consists of a Supermix® reactor. The reason that
this photochemical reactor was used in these studies is that mixing in the
reactor is experimentally efficient. There are two chambers in the reactor
which is equipped with a 85 ml recirculation loop, and has a working volume
of 250 m1. A quartz immersion well housing a Double-bore®

lamp was placed in the larger chamber in this reactor. A glass impeller shaft
(Model 8068-08, Ace Glass, Inc., Vineland, NJ) in the smaller chamber was

used to circulate the solution in the reactor. The shaft was rotated by

28

29

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

SAWUNE
GAS LINE
Wm
WM
pH ELECTRODE PORT
mam
-_ SHAFT
FLOW —— WWW
nmscnon LOOP
. l
* l _ mm was
s
(l k t
\
‘5-

Figure 3.1 Schematic of Supermix® Photochemical Reactor (Model #7868,
Ace Glass, Inc., Vineland, NJ)

30

using an electric motor (Model 102, Talboys Engineering Corp., Montrose,
PA).

3.1.2 Ultraviolet light source

The ultraviolet light source was manufactured by the Jelight Co., Laguna
Hills, CA. The lamp is a Double-bore® phosphor coated lamp (Model 84-2011-
8) with a peak output wavelength of 351 nm. It is a low pressure, cold cathode
phosphor coated mercury vapor lamp. The dimensions of the lamp are
shown in Figure 3.2. The lamp consists of quartz tubing with a quartz
partition or septum through the center. A Model PS-2000-20 power supply
(Ielight Co., Laguna Hills, CA) was used.

3.1.3 Experimental scheme

A schematic diagram of the apparatus used in this experiment is shown in
Figure 3.3. There are three sample ports at the top of the reactor. All the tubes
connected to the reactor were made of Teflon®, except the tubings used for

cooling loop.

To maintain a constant temperature in the reactor a peristaltic pump (Model
503U, Watson-Marlow, Concord, MA) was used to cycle cooling water

through the cooling jacket of the reactor.

31

M

 

 

 

 

 

j /- 0.125 no: I
-9.“ - ..... .

i 4+ l-m i

i LIGHTED LENGTH ~— 200 -—~ 3

0 o’

 

 

OPTIONAL INDEX SLOT

  

Figure 3.2 Structure of UV lamp

32

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

_/ \
\_ _
1
UV Lamp
Temperature —>
Control . ’ UV Lam
... A v
ys ’ Power Supply
Peristaltic
Pump .
Supermix® photoreactor

Figure 3.3 Experimental Scheme

33

To measure the pH in the solution in the reactor a pH electrode was located at
the sampling port. To maintain the required pH in the reactor sodium
hydroxide or sulfuric acid were, when necessary, introduced through tubing
into the bottom of the smaller chamber of the reactor.

Air or pure oxygen were bubbled into the bottom of the smaller chamber
through Teflon® tubing.

3.1.4 pH control system

A pH meter (Model 720 Laboratory Instrument, Orion Research Incorp.,
Boston, MA) connected to a pH probe ( Model 91-16 Orion Research Incorp.,
Boston, MA) a semi-micro pH electrode with a length of 180 mm and body

diameter of 6 mm were used to measure pH in the reactor.

3.1.5 High Performance Liquid Chromatograph

A Gilson high performance liquid chromatograph (HPLC) (Middleton, WI)
equipped with two model 303 piston pumps, a 8028 manometric module, 811
dynamic mixer, 116 UV detector, 231 auto-sampling injector, and 401 dilutor
was used for the 2-chlorophenol analyses. A C-18 PR column (Whatman,
Clifton, NJ) was used for the chromatographic separations. The C-18 column
has a length of 235 mm, CD. of 7.94 mm, ID. of 4.70 mm, void volume of 2.38
ml, and the particle size of the packing was 5 um. A mobile phase consisting
50% acetonitrile and 50% of 1% phosphorus acid in water was used. The flow

rate was 1.5 ml / min and a sample volume of 20 )1] was used. The UV detector

34

was set at 230 nm. The retention time for the peak of 2-chlorophenol was 3.6

min.

All the samples containing titanium dioxide will be filtrated through 0.2 um
or 0.22 um filters before analyzing by HPLC.

3.1.6 Spectrophotometer

A UV-visible recording spectrophotometer (Model UV 160, Shimadzu corp.,
Kyoto, Japan) was used to measure the absorbance of the filtering solution
and for the actinometry. 1 cm pathlength cells (from Spectrocell corp.,

Oreland, PA.) were used for these determinations.

3.1.7 Sample collection

A sampling tube located at the bottom of the smaller chamber in the
Supermix® reactor was used for sample collection. A 15 to 20 ml sample was
collected and filtered through 0.2 um (Scientific resources Inc., North
Brunswick, N]., and Gelman Science, Ann Arbor, MI.) or 0.22 um pore size
filter (Millipore Products Division, Bedford, MA.).

3.1.8 Nonlinear fit computer program

Systat ver. 5.2.1 for Macintosh was used as the nonlinear fit program in this

study.

35

3.2 Experiments:
3.2.1 Materials

Acetonitrile and ferrous chloride were purchased from EM Science,
Gibbstown, NI. Ferric sulfate, hydrogen peroxide, and 2-chlorophenol were
purchased from Sigma Chemical Co., St. Louis, MO. Sodium bicarbonate,
sodium hydroxide, sulfuric acid, and pH 7 calibration buffer solution were
obtained from ].T. Baker Inc., Phillipsburg NJ. Methanol was purchased from
Baxter Healthcare Corp., Muskegon, MI. Sodium carbonate was purchased
from Mallinckrodt Inc., Paris, Kentucky. 1,10—phenanthroline was purchased
from Aldrich Chem. Co., Milwaukee, WI. Phosphoric acid, and potassium
dichromate were obtained from Mallinckrodt Chemical Works, St. Louis,
MO. pH 4 calibration buffer solution was purchased from Baxter Healthcare
Corp., McGraw Park, IL. The photocatalyst Titanium Dioxide P25 was obtained
from Degussa Corp., Ridgefield Park, N]. It consists mainly of anatase. It has
an average primary particle size 30 nm, surface area 50 m2/ g, and PHzpc 6.3.
P25 contains some impurities such as A1203 (<0.3%), Si02 (<0.2%), Fe203
((0.01%), and HCl (<0.3%).

3.2.2 Chemical preparations

Chemical and solvents were of reagent grade and used without further
purification. Stock solutions of 2-chlorophenol (10000 ppm), sodium
bicarbonate (6.875x10'3 M), ferrous chloride (5.5x10‘3 M), potassium
dichromate (1x10'7- M), ferric sulfate (2x10’1 M), I<2C204 (1.2 M), sulfuric acid

(1 x1 0'1 M), and sodium hydroxide (1 x10‘2 M), were prepared and used in the

36

same day. All solutions were prepared by dissolving the appropriate amount

of chemical in deionized water.

3.2.3 Standard calibration curve

Five 2-chlorophenol standards (1 to 10 ppm) were used to prepare calibration
curve for HPLC. A calibration curve was prepared for each experiment by
plotting peak area vs. 2-chlorophenol concentration.

3.2.4 pH calibration

The pH meter was calibrated using two buffer solutions.

3.2.5 Batch experiments-UV/Titanium dioxide system

The experimental conditions used in the experiments were as follows:

Volume of solution 275 ml

Ti02 concentration 1 g/ L
2-chlorophenol concentration 1 0 ppm
Gas flow rate 175 ml/ min

After the solution was placed in the reactor, recycling was started. Gas was
bubbled into reactor, when required. The suspension was stored one hour in
the dark before illumination to allow a sorption equilibrium between catalyst

and 2-chlorophenol to be achieved.

37

The pH was adjusted before the photolysis was commenced. Temperature was
controlled over the entire irradiation period by circulating cooling water to

maintain a temperature of approximately 300:1:2 K.

A light intensity of 3 mW/ cm2 was used during experiments unless

otherwise indicated.

3.2.6 Batch experiments-UV/Titanium dioxide/Hydrogen peroxide system

The reaction solutions were prepared as described in section 3.2.5. The desired
amount of hydrogen peroxide was added to the reactor immediately before

the UV light was turned on.

A series of experiments were conducted using different hydrogen peroxide
concentrations (0.001, 0.050, 0.0250 M) without Ti02 at equilibrium pH to
determine the effect of H202 on 2-chlorophenol degradation.

3.2.7 Effect of pH

A series of experiments at an initial pH of 3.6, 5, 6.5, 8, or 10 were conducted
using UV/ Titanium dioxide system and UV/ Titanium dioxide/ Hydrogen
peroxide system. In these experiments sulfuric acid or sodium hydroxide were

added to adjust the pH to the desired value.

38

3.2.8 Effect of light intensity

A series of experiments at various light intensities (100%, 70%, 58%, or 35%
lamp output) were conducted using UV/ Titanium dioxide system and

UV/ Titanium dioxide/ Hydrogen peroxide system. In these experiments a
potassium dichromate solution containing 3% sodium carbonate was used as
a filter solution to decrease the light intensity. This solution was cycled

through the cooling jacket in the photoreactor.
3.2.9 Potassium ferrioxalate actinometry

The potassium ferrioxalate actinometer developed by Parker and Hatchard
and Parker (33) was used to measure light intensity. Fe2+ was determined by

the phenanthroline method (34). The net reaction is the following (35):
217a3+ + C2042' —"'—-)2Fe2+ + 2602 (44)

The Fe2+ is subsequently determined via spectrophotometric determination
of its phenanthroline complex at 51003. Fe3+ forms only a weak complex

with phenanthroline which is transparent at 5100/01.
3.2.10 Effect of illuminated method

Two different illumination methods were used. One is referred to as the

continuous method, the other is referred to as the intermittent method.

39

The continuous method was used unless otherwise indicated. In these
experiments the UV light was switched on, and was not turned off until the

end of this experiment.

For the intermittent method the UV power was turned on for 2 minutes then

turned off for 3 minutes. Sample collection was during the dark period.

3.2.11 Effect of oxygen concentration

A series of experiments were conducted by bubbling air or pure oxygen into
the reactor containing titanium dioxide or titanium dioxide with hydrogen

peroxide. Experiments were also conducted without any bubbling.

3.2.12 Effect of Fe2+

A series of experiments were conducted in 1x10”4 M ferrous chloride. Ferrous
chloride was put into reactor at least one hour before turning on the UV light

or adding hydrogen peroxide.

3.2.13 Buffer system

A series of experiments were conducted using different concentrations of
sodium bicarbonate (0.5 and 2.5 mM) to find the relationship between the rate
of decomposition of 2-chlorophenol and concentration of sodium bicarbonate

in solution.

Chapter 4

Results and Discussion

The rate of the photocatalytic degradation of 2-chlorophenol depends upon
various parameters including initial 2-chlorophenol concentration, pH, light
intensity, temperature, , wavelength of incident light, mass and type of

photocatalyst, and the design of photoreactor.

In this work the effect of pH, light intensity, concentration of hydrogen
peroxide, oxygen, and ferrous ion on the rate of photocatalytic degradation of
2-chlorophenol was studied. The initial conditions used in the experiments

conducted are summarized in Table 4.1.

40

Nt—i

13

14

15

16

41

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

T 14.1Thini'l ni'o ofx rimen
No. Experimental systems Illuminated pH Other control
Methodal Controlb parameters
UV / Airc Continuous None
UV/H202 Continuous None H202=0.001,
0.050, 0.250 M
UV [11022 Continuous None
UV /T102/ Air Continuous None
UV [1102 /()2f Continuous None
UV / Ti02/ H202 Continuous None H202=0.001 M
UV/Ti02/02 Intermittent 7.58
UV/TiOz/H202 Intermittent 7.19 H252=o.oo1 M
UV / Ti02 Intermittent 3, 5, 6.5, 8, 103
UV/ Ti02 Continuous 8.20 NaHC03=2.5 mM
7.51 NaHC03=0.5 mM
UV/Ti02/ Air Continuous 3.6, 5, 6.5, 8, 10 NaHC03=0.5 mM
UV/TiOz/Air/H202 Continuous 3.6, 5, 6.5, 8, 1o NaHC03=0.5 mM
H202=0.001 M
UV /Ti02 / Air/1232+}! Continuous 6.87 Marie—03:05 mM
Fe2+=o.0001 M
UV / Ti02 / Air / Fe§+ / fintinuous 6.93 NaHC03=0.5 mM
H202 Fe2+=0.0001 M
H202=0.001 M
UV / Ti02 / Air Continuous NaHC03=0.5 mM
Light
intensity
100% 6.74
70% 6.84
58% 6.86
35% 6.92
UV/Ti02/Air/ H202 Continuous NaHC03=0.5 mM
Light H202=0.001 M
intensity
100% 6.85
70% 6.87
58% 6.86
35% 6.96

 

 

a: The illumination method is explained in section 3.2.11.

 

42

b: In this study pH was adjusted to desired value prior to illumination.

c: Flowrate of gas was 175 ml/ min.

d: None means pH was not adjusted.

e: Concentration of Ti02 was 1 g/l.

f: Flowrate of pure oxygen is 175 ml/ min.

g: The pH was maintained at constant value in these experiments. In other
experimental systems the reported value is the initial pH.

h: Initial concentration of Fe2+ is 0.1 mM.

The results of the experiments are described in following sections. The

experimental data are given in Appendix A.

4.1 UV combined with air, H202, or Ti02

4.1.1 UV/Air

Because air was bubbled into system, an experiment was conducted to
determine how much 2-chlorophenol was lost due to volatilization.
Approximately 5.3% of 2-chlorophenol was lost only due to volatilization in a
two hour experiment (including in dark period). The dissolved oxygen

concentration with bubbling air in the solution was 8.88 mg/l.

In an illuminated solution, the rate of 2-chlorophenol disappearance was
slow. The loss of 2-chlorophenol after 60 minutes was approximately 10% (see
Table A1). If volatilization is considered, the net disappearance of 2-

chlorophenol due to direct photolysis was about 4.7%. The pseudo first order

43

rate constant including loss due to volatilization for the disappearance of 2-

chlorophenol was 0.002 1 / min.

The reason for bubbling air is to more closely simulate the conditions used in

photocatalytic degradation experiments.

4.1.2 UV/HzOz

The effect of initial concentration of hydrogen peroxide on the degradation

rate is shown in Figure A.2.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

UV/ H O
2 2

a 1 00
E 80
‘2 .
'5 60
o
g: .
= 4 0 ~—-
g _ O 0.001 M hydrogen peroxide
Q 2 _ a 0.050 M hydrogen peroxide

0 0 0.250 M hydrogen peroxide

0 r l r l m l 1 l 1 l

0 1 0 2 0 3 0 4 O 5 O 6 0

Time (min)

Figure A.2 2-chlorophenol decomposition in UV/H202 oxidation system

At an initial hydrogen peroxide concentration of 0.001, 0.050, and 0.250 M , 2-
chlorophenol was decomposed 9%, 24%, and 45%, respectively, after 60 min

44

illumination. The rate constant for the disappearance of 2-chlorophenol were

0.001, 0.005, and 0.010 1/ min, respectively.

Because the UV produced by the lamp used in this study is not energetic
enough to decompose hydrogen peroxide to -OH radicals directly, the rate of 2-

chlorophenol degradation in hydrogen peroxide solution is slow.

At 0.001 M hydrogen peroxide which was the concentration of hydrogen
peroxide used in subsequent experiments, the rate of 2-chlorophenol
degradation in hydrogen peroxide was not‘ significantly different than that

without hydrogen peroxide present.

4.1.3 UV/TiOz

Figure A.3 shows the photodegradation for 2-chlorophenol in an UV/ Ti02
system. 96% of 2-chlorophenol was degraded within 60 min. and the t1 /2 for
the degradation of 2-chlorophenol was 6.0 min. The first order rate constant

for the degradation of 2-chlorophenol was 0.115 1 / min.

As can be seen in Figure A3 at time greater than 20 minute, the pseudo-first
order kinetics does not fit the experimental data very well. This may be due to
a decrease in the oxygen content of the suspension. Since if there is not
enough oxygen to trap electrons excited into the conduction band then the
accumulation of electrons in the Ti02 particles will increase the
recombination rate of electrons and holes and thus decrease the

photodegradation rate.

45

 

 

 

 

 

 

 

 

 

 

UV/Tio2
100
g
i 80
a? _
3 60
8t .
'3 40
g .
20
0 b 9 ? o 9
0 1o 20 30 4o 50 60

Time (min)

Figure A.3 Photodegradation of 2-chlorophenol in UV / Ti02 system.

The stoichiometry of the reaction is the following.

C6H40HC1 + 5.502 —> 6C02 + 5H” + Cl’ (45)

Based on this stoichiometry one would expect in closed system 10 ppm of 2-
chlorophenol would consume about 13.7 mg/l of oxygen. The initial
concentration of oxygen in this study was 7.5 mg/l, so significant oxygen
depletion will occur during the experiment. Thus in future experiments
bubbling oxygen or air into the system was used to control the oxygen level in

the suspension.

46

4.1.4 Comparison of UV radiation combined with air, H202, or Ti02

From the experimental results shown in Figures A.1, A2 and A.3, the effect of
UV radiation, H202, and Ti02 can be determined. It is concluded that

The direct photolysis of 2-chlorophenol is very slow (see Figure A.1).

0 The degradation rate increases with increasing hydrogen peroxide
concentration, but the degradation rate is not proportional to the

concentration of hydrogen peroxide.

0 The photodegradation rate of 2-chlorophenol in the photocatalytic system

is much faster (about 60 times) than that the rate of direct photolysis.

0 The 2-chlorophenol degradation rates in the UV-Air and UV-0.001 M
H202 are almost equal, i.e., the rate of reaction of 2-chlorophenol with

H202 is not significant at low H202 concentration.

4.2 UV/TiOz combined with air or pure oxygen

4.2.1 UV/TiOz/Air

In this experiment, air was bubbled into the reactor to maintain a constant

oxygen concentration in UV / Ti02 system. The result of this experiment is

shown in Figure A.4.

47

99% of 2-chlorophenol was degraded during 35 min. and the t1 /2 was 5.5

min. The rate constant was 0.125 1/ min.

UV/Ti02/Ai r

 

 

100

 

80
so.
a)
2. \

0P.1%N.1.1

0 10 20 30 4O 50 60

Time (min)

 

 

Percentage of Remaining

 

 

 

 

Figure A.4 Photodegradation of 2-chlorophenol in UV / Ti02 / Air system

Bubbling air into this system increased the reaction rate slightly suggesting
that in the experiment described in the previous section oxygen depletion was
occurring. No significant deviation from first order kinetics was observed in

the later stages of this experiment.

4.2.2 UV/TiOz/Oz

In this experiment, pure oxygen was bubbled into the UV / Ti02 system to
examine the effect of oxygen concentration on the reaction kinetics. The

result of this experiment is shown in Figure A.5.

48

97% of 2-chlorophenol was degraded during 20 min. and the t1 /2 was 4.8

min. The rate constant was 0.145 1/ min.

UV/Ti02/02

 

 

100
80

.\
60 K
40

\
\\

OPIIIQIII l 1L4
o 10 20 so 40 50 so

Time (min)

 

 

 

Percentage of Remaining

 

 

 

 

Figure A.5 Photodegradation of 2-chlorophenol in UV/ Ti02/ 02 system.

Bubbling pure oxygen into this system instead of air increased the reaction
rate by 16%. The oxygen concentration in the solution with bubbling pure

oxygen was 16.48 mg/l.

4.2.3 Effect of 02 concentration on UV/TiOz, UV/TiOz/Air, and UV/TiOz/Oz

system

A comparison of experimental results for the decomposition of 2-

chlorophenol in UV / Ti02, UV/TiOz/Air, and UV/ Ti02/ 02 is presented in

Figure 4.1.

49

Effect of O2 Concentration

 

 

 

 

 

 

 

 

    

 

0.20
A El UV/TiO
E UV/Tl02/Air
E 015 _... B UV/'l"102/O2
E
2
2
8 010 — --------
2
S
t:
.2
‘5 0.05 - """""
3
I

0

 

 

 

Figure 4.1 Effect of 02 concentration on photodegradation of 2-chlorophenol
in UV / Ti02, UV / TiO2/ Air, and UV/TiOz/Oz system.

Our experiments show that while bubbling oxygen or air into the system does

increase the rate of 2—chlorophenol degradation, but the extent of the increase

is small.

From these studies on the effect of oxygen concentration of the degradation of

2-chlorophenol, the following conclusions can be made:

0 Bubbling air into UV/Ti02 system increased reaction rate by 9%.

0 Bubbling pure oxygen, which increased the oxygen concentration in

solution about two times compared to that when air was used, increased

50

photocatalytic degradation rate by 26%. This shows that the reaction rate is

not proportional to the concentration of oxygen in solution.

0 With bubbling air or pure oxygen into the UV / Ti02 system it was found

that the first order kinetics can fit the data very well.

Al-Ekabi et al. (19) have found similar results to those obtained in our
experiments. Their data indicated that oxygen plays a crucial role in the
photocatalytic degradation of 2,4-dichlorophenol. They mentioned that
introducing small bubbles of either air or oxygen may improve the mixing

process inside the reactor and thus increase the reaction rate.

The adsorption of oxygen can be described by a simple Langmuir adsorption

isotherm as follows:

—-—2—L—K [O] (46)

x°= -1+Ko[02]

where 102 is the fraction of total surface sites occupied by 02 ([0215urf= lo, [02]
and K0 is the surface binding constant for 02 on Degussa P25 Ti02. K0 was
found to be 132t7x104 M‘1 (22).

From our results, 101 were 0.973 and 0.985 for bubbling air and pure oxygen,

respectively. These indicate that the surface sites were almost occupied by

oxygen.

In our experiments, we assume that conduction band electrons are scavenged

by 02 or by an alternate electron acceptor (e.g., H202) while htm has a slightly

51

longer lifetime with respect to electron transfer (i.e., fast recombination is no
longer possible). During this longer lifetime, the holes are scavenged by the
hydroxylated surface of Ti02 to produce OH radicals (36).

Figure 4.2 shows the photocatalytic degradation rate of 2-chlorophenol is a
nonlinear function of oxygen concentration in solution. A simple Langmuir

adsorption can be used to fit the experimental data (r2=0.995).

Effect of [02] on UVI'I‘iO2 system

 

 

 

 

 

 

 

 

 

 

 

 

 

0.2
A 0 experimental data
E Langmuir adsorption m
g f .
‘6
S
o 0.1
3
t!
C
.2
‘5 0.05
a
e
(I

o c l I i I i I

1 0 1 5 20
[02] (mg/l) in solution

Figure 4.2 Effect of oxygen concentration on the photocatalytic degradation

rate of 2-chlorophenol.

52

4.3 UV/TiOz combined with H202

4.3.1 UV/TiOz/HzOz

In this experiment 0.001 M of hydrogen peroxide was added to the UV/Ti02

system. The result of this experiment is shown in Figure A.6.

99% of 2-chlorophenol was degraded during 20 min. and the t1 / 2 was 2.9 min.

The rate constant was 0.235 1 / min.

The reason why the reaction rate is so high is that hydrogen peroxide may
trap electrons in the conduction band (see reaction 23 in section 2.2). It also

produces oxygen which reacts with conduction band electrons.

Furthermore hydrogen peroxide can react with holes generated by titanium

dioxide to produce H02- which can react with 2-chlorophenol and degrade it.

The results of this experiment are in agreement with those of Auguliaro et al.
(30) which showed that the rate of photocatalytic degradation of phenol was
increased in the presence of hydrogen peroxide. Tanaka et al. (28) have also

shown that the addition of hydrogen peroxide to a Ti02 suspension resulted

in an increase in the rate of photocatalytic degradation of TCE.

4.3.2 Comparison of UV/I-I202 and UV/TiOz/HzOz system

From Figure A2 and Figure A.6 the reaction rate in UV/Ti02/ H202 system is
about 200 times faster than that in UV/H202 system. It can be concluded that

53

the addition of hydrogen peroxide to a Ti02 suspension can improve the

efficiency of the photocatalytic degradation of pollutants.

4.4 Photocatalytic systems under intermittent illumination method

4.4.1 UV/TiOz/Oz system under intermittent method

The use of utilizing intermittent illumination was investigated as the
reaction time of continuous method in UV/ Ti02 / Air, UV/Ti02/02, or
UV/Ti02/ H202 system was too short to allow time to collect and process
samples. By using the intermittent method it was possible to collect samples
during the period when the sample was not illuminated. The results of an
experiment conducted using intermittent illumination are present in Figure

A.7.

81% of 2-chlorophenol was degraded during 26 minutes of illumination time
and the t1 / 2 was 12.4 min. The rate constant for the degradation of 2-
chlorophenol was 0.056 1 / min which is 160% less than that obtained using

the continuous method of illumination (see Figure 4.3).

Furthermore, there was considerable scatter in the experimental data (see

Figure A.7).

54

4.4.2 UV/TiOz/HzOz under intermittent method

In this experiment the effect of hydrogen peroxide in UV/Ti02 system under
intermittent method was studied. Figure A.8 shows the degradation of 2-

chlorophenol during this experiment.

After 26 minutes of illumination, 85% of 2-chlorophenol was degraded, and
the t1 / 2 for the disappearance of 2-chlorophenol was 8.1 min. The rate

constant for the degradation of 2-chlorophenol was 0.086 1 /min.

The results indicate that the rate constant is 170% smaller than that in the
same system using continuous method (see Figure 4.3). Also there was

considerable scatter in the experimental data (see Figure A.8).

4.4.3 Effect of intermittent method on UV/TiOz/Oz and UV/TiOz/HzOz

system

The results of the photodegradation of 2-chlorophenol in UV/ Ti02/ 02 and
UV / Ti02/ H202 system under intermittent method show that

0 The reaction rate is slower under intermittent illumination than that
under continuous illumination. This suggests that the UV lamp needs

time to reach full power after it is turned on.

0 Under intermittent illumination the photodegradation is not well
described as first order kinetics. This is consistent with the fact that the

light intensity varies during the time which UV lamp is turned on.

55

A comparison of the photocatalytic degradation of 2—chlorophenol in
UV/TiOz/H202 and UV/TiO2/02 system under different illumination

method is shown in Figure 4.3.

Comparison of illumination methods

 

 

 

 

 

   

  

0.25 _
A : L D Intermittent method
C )- ’55
E 0'20 Tn" I Continuous method ...........
E I
a 0.15 g
C
o _
0 r—
o .—
at 0.10 _— ........................................................
c -
.9 -
8 -
g 0.05 f ...........
I —
0

 

 

 

 

 

uvmoz/oz UV/TiOZ/HZOZ

Figure 4.3 Comparison of different illumination methods in same oxidation

system.

These experiments indicate that intermittent illumination is unsuitable as a

method to slow down the reaction rate.
4.5 UV/TiOz under intermittent method and pH control

The other purpose of utilizing intermittent method is that it is easier than

continuous method to control pH at certain level during photocatalytic

56

degradation. When the UV lamp was off acidic and basic solution were added

to maintain certain pH value in solution.

In this experiment pH 3, 5, 6.5, 8, and 10 were controlled to perform this
photooxidation reaction. The results of these experiments are shown in

Figure A.9, A.10, A.11, A.12, and A.13. The results are present in Table 4.2.

Tal 4.2Ther l f2-chlr hnl e 'nin VTi_ stem
gnder intermittent method at differgnt pH during 26 min illumination

—_-. rate constant u/min)
3 6.4

84%

   
    
   
   
   

 

    

 

90% 8.5

 

86%

 

86%

 

52%

 

 

 

The fitting results were not good due to intermittent illumination, but

constant pH was maintained during the experiment. Because intermittent
illumination is not a good method to use, for other experiments pH was not

held at a constant value.

4.6 UV/TiOz under buffer system

57

Since intermittent illumination method cannot be used to slow down the
degradation rate, an alternative method to slow down the reaction is to use a

bicarbonate buffer.

To determine the effect of bicarbonate ion concentration on the reaction
kinetics experiments, a number of experiments were conducted in 2.5 and 0.5

mM sodium bicarbonate solution.

As seen in Figure A.14 and A.15, 74% and 92% of 2-chlorophenol was
degraded during 60 min in 2.5 mM and 0.5 mM NaHC03 solution,
respectively. The t1 / 2 for the disappearance of 2-chlorophenol was 28.9 and
11.4 min, and the rate constant for the degradation of 2-chlorophenol was

0.024, and 0.061 1/ min, respectively.

These experiments indicate that

0 The higher the concentration of sodium bicarbonate the lower the reaction
rate. HCO3' is a -0H radical scavenger, and the rate constant for the
reaction between HCO3' and OH is 8.5x106 M‘IS'1 (37). The addition of
bicarbonate could slow down the reaction as ~0H radicals generated are

scavenged by HCO3'.

0 All subsequent experiments were conducted with 0.5 mM NaHCO3 buffer
solutions as this concentration is sufficient to slow down the reaction to

the desired degree.

UV/T iOzl NaH C 03 photocatalytic system

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

0.12
i :1 without iii-Moo
g 010 * ,,,,, I witito.5mMNaH003 ,,,,,,
g ' . :1 with 2.5 mliiiNaHoo3
E 0.08
2
a .
g .-
o 0.06
2 I
g _ -.-..--
e 0.04 1:: E
8 i .1___.
E 0.02 _ /
0 fl

 

 

 

 

 

 

Figure 4.4 The effect of HC03' on the photocatalytic degradation rate of 2-

chlorophenol.

4.7 Effect of pH on photocatalytic systems

4.7.1 UV/TiOz/Air/NaHCO3 with and without H202 under pH control

In these experiments the initial pH was adjusted to 3.6, 5, 6.5, 8, and 10. The
results at different initial pH value with and without hydrogen peroxide are
shown in Figure A.16, A.17, A.18, A.19, and A20. The results are summarized

in Table 4.3. The effect of initial pH on rate constants is shown in Figure 4.5.

59

Tl4.Thr l f-hlrhnl r 'nin
UVZTinJAirZNaHQQg with ma withgut H_2Qz_at different initial pH

UV/TiOz/Air/Nal-ICOa with H202
m rate constant tt/ntn)
3.6 6.5

5 99% 6.4 0.108

         

  

 

   
    
    
 
 
   

 

 

8.2
8.9

UV/TiOz/Air/NaHCO3 without H202
—_- constnn tt/ntni

 

 

 

 

 

  

 

 

 

 

 

 

 

60

Effect of pH on UV/TiOzlAir/NaHCO3

 

 

  
 
 

 

  

 

 

with and without H202
0.1 2
' K 0 without hydrogen peroxide
E 0.1 0 ~ """ L """ with 0.001 M hydrogen peroxide “
E 0.08 . .......... g ........................
+3
§ 0.06 . ........................
3
S
c 0.04 . ....................
.9 . :
g . i.
(E 0.02
0
3.5 5 6.5 8 1 0
Initial pH

Figure 4.5 Effect of initial pH on UV/TiOz/Air/NaHCO3 system with and
without H202.

The results show that without hydrogen peroxide the highest reaction rate

was observed when the initial pH was 6.5.

Similarly, the result also shows that when hydrogen peroxide was added the
highest reaction rate was observed in the experiment when the initial pH was

5.

4.7.2 Effect of pH on UV/TiOz/Air/NaHCOa with and without H202

Before discussing the effect of pH on the photocatalytic degradation rate of 2-

chlorophenol, some information needs to be reviewed. The pKal and pKaz of

H2C03 at 25 °C are 6.3 and 10.3, respectively. The rate constants from Buxton

61

et al. for the reaction of OH and HC03’, and for the reaction of -OH and C032‘
are 8.5x106 M"ls‘1 and 3.9x108 M454, respectively (37).

The following observations on the effect of pH on the photocatalytic
degradation of 2-chlorophenol can be made:

0 The difference of rate constants decreased from pH 3.5 to 10. The decrease
in the rate of photocatalytic degradation of 2-chlorophenol with pH can be
attributed to the scavenging of -OH radicals by bicarbonate and the rate of
reaction (~0H radicals with carbonate) which is about 46 times higher than
that with bicarbonate. At high pH the photocatalytic degradation rate is

much slower.

0 At lower pH such as pH 3.5 the photodegradation was slow without
hydrogen peroxide. The decrease in the rate seen at low pH can be
attributed to low concentration of OH' ions, because OH' ions react with
holes to generate -OH radicals. D’Oliveira et al. found the similar decrease

in the rate of photocatalytic degradation of 2- and 3-chlorophenol in Ti02

aqueous suspensions (23).

0 From Figure 4.5 at an initial pH in the range 5-8, no significant change of

photocatalytic degradation rate of 2-chlorophenol was observed.

62

4.8 Effect of Fe2+ on photocatalytic systems
4.8.1 UV/TiOz/Air/NaHCO3/Fe21’ with and without H202

In this experiment the effect of the addition of ferrous chloride on UV / Ti02
system was studied. The results of this experiment at equilibrium pH (6.9)
with and without hydrogen peroxide are shown in Figure A.21.

Without hydrogen peroxide addition 87% of 2-chlorophenol was degraded in
35 min., the t1 / 2 for the disappearance of 2-chlorophenol was 15.8 min. , and
the rate constant was 0.044 1/ min. With 0.001 M hydrogen peroxide addition,
91% of 2-chlorophenol was degraded, the t1 /2 for the disappearance of 2-
chlorophenol was 10.8 min. , and the rate constant for the degradation of 2-

chlorophenol was 0.064 1/ min.

The reason to add ferrous ion to the system is that in this experiment -OH
radicals could be generated from photooxidation, and produced through
chain reactions by Fenton or Fenton-like reactions. Hydrogen peroxide in this
system could be generated as result of the photocatalytic process (see reactions

in chapter 2).
4.8.2 Effect of adding ferrous chloride
The comparison of the rate constants for the photocatalytic degradation of 2-

chlorophenol with and without ferrous chloride addition is shown in Figure

4.6. The following observations can be made:

63

0 The reaction rate constants in UV/TiOz/Air/NaHCOs/Fe2+ system are
smaller than those in UV/TiO2/Air/NaHC03 system.

0 The addition of hydrogen peroxide enhances the photocatalytic
degradation rate.

0 The addition of ferrous chloride did not increase the degradation rate.

Effect of 1=e2+ on UV/TiOzlAir/NaHCOS

U lN/TDZIAlr/NaHooalFez'
n LN/TDZINr/Natma

 

0.15

 

0.05

Reaction rate constant (1/min)

 

 

 

 

 

 

with H202 without H202

Figure 4.6 Effect of Fe2+ on UV/TiOz/Air/NaHCO3 with and without H202,

64

4.9 Effect of light intensity on photocatalytic systems

4.9.1 UV/TiOz/Air/NaI-ICO3 with and without H202 under different light
intensity

To study the effect of light intensity on the reaction kinetics a number of
experiments were conducted under different light conditions. The results of
these experiments are shown in Figure A22, A23, A24, and A25. These

results are summarized in Table 4.4, and Figure 4.7.

_T_able 4.4 The results of 2-chlorophenol degradation in
UVZTin_ZAirZNaH§Q_3 with and wiQQut 202 at diffgrent light intansity

UV/‘I‘iOz/Air/NaHCO3 with H202
. I
111111ng _ t1/2 __ 1 rate Wst <1/mi__

100% 99%

 
 
  

 
  
  
   

    
 

 

     
    
  

70% 96%

 

58% 92%

   

 

  
     

 

 

 

UV/TiOz/Air/NaHCO3 without H202
? 1' |
W“ __ no

737

 

 

 

 

 

 

 

 

65

Effect of lighty intensity on UV/TiOzlAir/NaHCOS

 

0.15

D without hydrogen peroxide
' B with 0.001 M hydrogen peroxide

Ii ~ 3'.

1 00% 70% 58% 35%
Light intensity

Figure 4.7 Effect of light intensity on UV/TiOz/Air/NaHCOa system with and

 

0.10

0.05

Reaction rate constant (1/min)

        

 

 

   

without H202,

4.9.2 Effect of light intensity

The following observations can be made of the influence of light intensity on

the reaction kinetics.

0 The higher the light intensity the higher the reaction rate.

0 The difference between the system with and without adding hydrogen

peroxide of rate constants decreased as light intensity decreased.

66

0 At 100% of light intensity the effect of adding hydrogen peroxide was
significant. However, at the 35% of light intensity the addition of

hydrogen peroxide did not significantly increase the reaction rate.

0 At higher light intensity more electrons and holes were generated on the
surface of titanium dioxide. Hydrogen peroxide can react with those
electrons to generate oOH radicals. However, at lower light intensity, fewer
electrons and holes were formed on the surface of titanium dioxide. The
effect of the addition of hydrogen peroxide was very small at 35% of light

intensity.

Figure 4.8 shows without the addition of hydrogen peroxide the rate of 2-
chlorophenol photodegradation is a nonlinear function of light intensity.
However, when hydrogen peroxide was added the photocatalytic degradation
rate is a linear function of light intensity (r2=0.997). A square root function
can be used to fit the data from the experiments without hydrogen peroxide
addition (r2=0.999).

67

 

 

 

 

 

 

 

 

 

Effect of light intensity
on UV/TiOZIAir/NaHCO3 system
0.1 6
A 0 without H20 / 4L
E square root?“ . /
z: 0-12 P" A with H202 / ,
E - - - linear fit /,/
a ,- A
8 0.08 , it 2
3 ./
E L / O
c /
.2
g 0.04 /r
a? _ /'
/
o l/ l l g l L L 1 l 1

 

 

 

0 20% 40% 60% 80% 100%
Light intensity

Figure 4.8 Effect of light intensity on the rate of photocatalytic degradation of
2-chlorophenol with and without adding hydrogen peroxide addition (where
100%=3 mW/cmz).

Al-Sayyed et at. (18) and D’Oliveira et al. (23) have observed a similar
dependence of light intensity for the photocatalytic degradation of 4-

chlorophenol, and 2- and 3-chlorophenol, respectively.

The initial quantum yield for the photocatalytic degradation of 2-
chlorophenol without hydrogen peroxide addition at different light intensity
is shown in Table 4.5. These calculations assume a reflectivity of 0.15 (23). The
calculations shown are lower limits for the quantum yield, since no

allowance was made for scattering by the solid particles.

Tl4.Ini°1 n 'l -hlr henl

Liht intensi Initial cuantum ield

 

 

 

 

 

Al-Sayyed et al. (18) have present a initial quantum yield of 2-chlorophenol
was 0.013 and 0.009 at the 2 and 50 mW/cm2 light intensity, respectively.

D’Oliveira et al. (23) have shown a initial quantum yield of 2-chlorophenol
was 6.9x10‘3 at a radiant flux of 50 mW/cm2 absorbed by Ti02.

4.10 Summary of reaction rate constants

A summary of the experimental rate constants obtained in this study is

presented in Table A25.

Chapter 5

Conclusions and Recommendations

5.1 Conclusions

From the results of this study, the following conclusions can be made:

1. The reaction rate for photocatalytic degradation of 2-chlorophenol is fast.

2. The addition of hydrogen peroxide can enhance the photocatalytic
degradation rate of 2-chlorophenol significantly under appropriate
conditions. The enhancement in the rate in the presence of hydrogen

peroxide is greatest at low pH and under intense illumination.

3. Bubbling air or pure oxygen into UV/Ti02 system does not increase the rate

of degradation of 2-chlorophenol to a great extent.

69

70

4. The intermittent illumination method is not a good method to study the

photocatalytic reaction.

5. Bicarbonate and carbonate inhibit the rate of photocatalytic degradation of

2-chlorophenol.

6. The addition of ferrous chloride inhibits the photocatalytic degradation rate

of 2-chlorophenol.

7. The optimum pH range to operate the photocatalytic degradation without
adding hydrogen peroxide is between 5 and 8. With the addition of
hydrogen peroxide the optimum pH range for the photocatalytic
degradation of 2-chlorophenol is between 3.6 and 5.

5.2 Recommendations

In order to understand more detail about the mechanism, more experiments

should be considered in the further studies:

1. The effect of process parameters such as concentration of 2-chlorophenol,
concentration of Fe2+, concentration of reaction released products like
chloride, concentration of Ti02, other wavelength of source light and other
effects such as oxygen concentration on the reaction kinetics should be

studied.

2. The reaction pathway and intermediate products, by-products, and products
should be studied.

[1].

[2].

[3].

[4].

[5].

[6].

LIST OF REFERENCES

Al-Ekabi H. and Serpone N., ”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, 5726-5731 (1988).

Schraa G. and Zehnder A.].B., In: i Mi 11 t n Th

A 'Envir nmn:Pr ein Th 4hEuro ean rn im.
Bjorseth A. and Angeletti G., Eds. Riedel D. Publ., Dordrecht,
Netherlands. 278-291 (1986).

Sonzogni W.C. and Swan W.R., In: Toxic Contaminants In The Great
Lakes. John Wiley and Sons, New York, NY. 1-30 (1984).

Kool H.]., van Kreiji CF. and I-Irubec 1., In: Water Chlorination:

Chemisgy. Environmental Impact And Health Effects. Vol. 5. Iolley, R.
L., Bull R.]., Davis W.P., Katz 8., Robert Jr. M.H, and Jacobs V.A., Eds.

Lewis Publ., Chelsea, MI. 187-205 (1985).

Pearson, C. R., In: Handbook Of Environmental Chemistgy:
Anthropogenic Compounds. Vol. 3, Part B.O. Hutzinger, Ed. Springer-

Verlag, Berline. 89-116 (1982).

Krijgsheld K. R. and van der Gen A., ”Assessment Of The Impact Of The
Emission Of Certain Organochlorine Compounds On The Aquatic

71

[8].

[9].

[10].

[11].

[12].

72

Environment, Part I: Monochlorophenols And 2,4-dichlorophenol”,

Chemosphere, 15, 825-860 (1986).

Kirt-Othmer Encyclopedia of Chemical Technology, 3rd ed., Freiter, E.R.,
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APPENDIX A
Experimental data

Ta 1eA.1D fFi r A.1frme rimn stemN.1

Time Percentage Fi
Remainin

 

91

 

 

 

 

 

 

 

 

 

 

UV/Alr
E” 1001 o o o o 0 o
E so
E L-
"5 60
8.
9.
c 40
20
o p 1 l 1 l 1 l 1 I l I 1
0 10 20 30 40 50 60
Time(min)

Figure A.1 Photodegradation of 2-chlorophenol under UV light

77

T l A. D fFi r A.2frm x rim nt sstemNo.2

 

 

 

 

 

 

 

 

 

 

 

 

 

 

a
C
E
E
II
‘6
o
a t
g
g 4 o _. .........................................
o O 0.001 M hydrogen peroxide
°- 2 0 A 0.050 M hydrogen peroxide

. o 0.250 M hydrogen peroxide

0 | I I I 1 l 1 L 1 l 1
0 1 O 2 0 3 0 4 0 5 0 6 0

Time (min)

Figure A.2 2-chlorophenol decomposition in UV/HzOz oxidation system

Percentage
Remainin

l

 

UV/TiO2

 

 

100‘

 

 

 

Percentage of Remaining

 

 

 

‘3 9 ? 9
0 10 20 30 40 50 60

Time (min)

v

 

 

Figure A.3 Photodegradation of 2-chlorophenol in UV/TiOz system.

T 1A. fF' . m rimn m .4

Time (min) Percentage of Fitting value
Remainin

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

UV/TiOzlAir
i
100<
13’
'g 80
g _
'5 60
i: . \
c 40
g _ \
20
o P 1 l 1 ? (D 1 l 4 l I
0 10 20 30 40 50 60
Time (min)

Figure A.4 Photodegradation of 2-chlorophenol in UV / Ti02 / Air system

 

81

Tabla A: Data pf Figarg A5 frgm gmrimantal aystem No. 5

Time (min) Percentage of Fitting value
Remainin;

 

 

 

 

 

 

 

 

 

UV/Ti02/02

 

 

100

 

80

60 \i
40 \
20

 

 

Percentage of Remaining

 

 

 

 

o 1 l 1 (D 4 L 1 l 1 i 1
0 10 20 30 40 50 60
Time (min)

Figure A.5 Photodegradation of 2-chlorophenol in UV/TiOz/Oz system.

82

TbleA. D fFi r A. rimntl temN.6

Time (min) Percentage of Fitting value
Remainin;

 

 

 

 

 

 

 

UV/Ti02/ H202

 

 

 

 

 

 

100
3
§ 80
g .
'5 60
8. L\
2
g 40 \
a 1-
20

 

 

 

0 10 20 30 4O 50 60

Time (min)

Figure A.6 Photodegradation of 2-chlor0phenol in UV/Ti02/ H202 system

T 1A.7D fFi r .7frm ri n sem .7

Percentage
Remainin

 

UV/Ti02/O2 under intermittent method

 

 

1oo<

 

 

 

 

 

 

 

'g 80

g t \

”5 60 o

o 0 o

81 t \

g 40 e

a . K
20 ° C
0 h 1 L 1 1 1 L 1

0 5 10 15 20 25 30
Time (min)

Figure A.7 Photodegradation of 2-chlorophenol in UV/TiOz/ 02 system

under intermittent method

84

Table A.§ Data pf Figprg A.§ frpm amrimgntal syatem Na. 8

tage
Remainin

100.

 

UV/TiO2/H202 under intermittent method

 

 

100 ch
80 o

60 - \

4o 0

20 i \°

 

 

 

Percentage of Remaining
o
1/

 

 

 

 

\GQ
o 1 l 1 I 1 l 1 l 1 l 1
0 5 10 15 20 25 30
Time (min)

Figure A.8 Photodegradation of 2-chlorophenol in UV/TiOz/HzOz system

under intermittent method

85

Tal A. fFi r A. frm 'mn tmNo.

tage pH
Remainin

100.00 03

 

UV/TiO2 under intermittent method at pH=3

100
T\
80 \
60 o
i o\
40 o
O
20 "

0 5 10 15 20 25 30

Time (min)

 

 

 

 

 

Percentage of Remaining

 

 

 

 

Figure A.9 Photodegradation of 2-chlorophenol in UV / Ti02 system under

intermittent method at pH=3

86

Tabla 5.19 Data pf Figarg AJQ frgm samrimantal system No. 9

tage pH
Remainin

1

 

UV/TiO2 under intermittent method at pH=5

 

 

100<
80

60 \

40 90\
20 0 e

0 5 1O 15 20 25 30

Time (min)

 

 

 

Percentage of Remaining

 

 

 

 

Figure A.10 Photodegradation of 2-chlorophenol in UV / Ti02 system under

intermittent method at pH=5

87

T 1A.11D fFi A. frm ' n mNo.9

tage Fi
Remainin

1
91

 

UV/TiO2 under intermittent method at pH=6.5

 

 

1ooc

 

 

 

 

 

 

 

.3 0

E 30 \

é _ o

”5 60 o

q, o

g r \

S 40 0

a r \

20 O O \
i- O O
o J J l i l l
0 5 10 15 20 25 30
Time (min)

Figure A.11 Photodegradation of 2-chlorophenol in UV / Ti02 system under

intermittent method at pH=6.5

88

T 1A.1 fFi r .12frm rim mNo.9

tage Fi
Remainin

00.

 

UV/TiO2 under intermittent method at pH=8

 

 

100C 00

 

 

 

 

 

 

 

IE 0
E 80 _ \D
'5 60
o
3 ~ \
g 40 _ o o
20 ° ° 0 0‘
_ o
0 1 1 1 . 1 . 1 1 1 .
0 5 10 15 20 25 30
Time(min)

Figure A.12 Photodegradation of 2-chlorophenol in UV/TiOz system under

intermittent method at pH=8

'0‘ .1 L. -_f i ‘ . om‘su‘1_'1--

tage
Remainin

 

UV/TiO2 under intermittent method at pH=10

 

 

 

 

 

 

 

 

 

mm c
2’ . ° 0
:§ 0
g 80 9
a p \
.1 * .
40
g _
20
O 1 l 1 l L l J I 1 I
0 5 10 15 20 25 30

Time (min)
Figure A.13 Photodegradation of 2-chlorophenol in UV / Ti02 system under

intermittent method at pH=10

90

T l A.14D fFi A.14 fr m rim n tem o. 10

Percentage pH
Remainin

1
.1

 

UV/TiO2 under buffer system

 

 

100<

 

 

 

 

 

 

 

8
§ so a
11% ~ \
ff so 0
E P 0 O \
5 40 o o
o O O
at i o O >
20
L
0 I 1 J l 1 l L l 1
0 10 20 3O 4O 50 60
Time (min)

Figure A.14 Photodegradation of 2-chlorophenol in UV/TiOz system under
buffer system with 2.5 mM NaHCO3

91

T 1e A] D fFi A.1 fr m rim n tern No. 10

Percentage Fi
Remainin

 

UV/TiO2 under buffer system

 

 

100<

 

 

 

 

 

 

 

.3
'2 so -
a t\
'5 60
5: - °\
= 40 so
E - \.\.
20 ° 9
_ 0 o o
o 1 l 4 l 1 L 1 l 1 1
0 1O 20 30 40 50 6O
Time(min)

Figure A.15 Photodegradation of 2-chlorophenol in UV / Ti02 system under
buffer system with 0.5 mM NaHCO3

92

Tal A.16D fFi r A.1 x rimn em .11 11 No.12

 

UV/TiOZIAir/NaHCO3 with and without H202
at initial pH=3.6

0 without hydrogen peroxide
A with 0.001 M hydrogen peroxide

Percentage of Remaining

 

0 10 20 30 40 50 60

Time (min)
Figure A.16 Photodegradation of 2-chlorophenol in UV / Ti02/ Air/ NaHC03

system at initial pH=3.6

93

Tbl A.17D fFi r A.17 x rimn mN .11 No. 12

 

UV/TiOZ/Air/NaHCO3 with and without H202
at initial pH=5

0 without hydrogen peroxide

1 00 A with 0.001 M hydrogen peroxide

80
60

40

Percentage of Remaining

20

O

 

0 10 20 30 40 50 60

Time (min)
Figure A.17 Photodegradation of 2-chlorophenol in UV / Ti02 /Air / N aHCO3

system at initial pH=5

94

T l A.18D fFi r A.1 ex rimen t mNo.11 No. 12

 

UV/TiOZ/Air/NaHCO3 with and without HAO2
at initial pH=6.5

0 without hydrogen peroxide
A with 0.001 M hydrogen peroxide

100

80

60

40

Percentage of Remaining

20

 

0 10 20 30 40 50 60

Time (min)
Figure A.18 Photodegradation of 2-chlorophenol in UV / Ti02 / Air/ NaHCO3

system at initial pH=6.5

95

Table A.12 Data 9f Figars A.12 exmrimentg systgm N9. 11 and No. 12

 

UV/Ti02/Air/NaH003 with and without F502

at initial pH=8
0 without hydrogen peroxide
2 1 00 A with 0.001 M hydrogen peroxide
‘6
'3 so
1:"
'5 6 0
8.
g 4
g 0
a:
20
0

 

0 10 20 30 40 50 60

Time (min)
Figure A.19 Photodegradation of 2-chlorophenol in UV / T102 / Air/ NaHCO3

system at initial pH=8

96

T leA20D fFi r A.2 x rimn mNo.11 N.12

 

UV/Ti02/Air/NaHCOa with and without l-lZO2
atiMfim pH=10

 

 

0 without hrdrogen peroxide
A with 0.001 M hydrogen peroxide

 

 

 

 

 

Percentage of Remaining
A O)
O O
o

I)
/
/

o

 

 

 

 

0 10 20 30 40 50 60
Time (min)

Figure A.20 Photodegradation of 2-chlorophenol in UV/TiOz/Air/NaHCO3

system at initial pH=10

97

T 1A.21Da fFi r A.21frm 'mnl mNo.1 dNo.14

 

UV/Ti02/Air/NaHCOS/Fez“ with and without H202

o wiyhout hydrogen peroxide
A with 0.001 M hydrogen peroxide

Percentage of Remaining

 

0 510 152025 303540

Time (min)
Figure A.21 Photodegradation of 2-chlorophenol in UV/ TiOz/ Air / Fe2+

under buffer system

98

Tabla A.22 Data at Figars A.22 frgm gmrimgntal systsm No. 15 and No. 16

 

UV/Ti02/Air/NaHCO3 with and without l-lZO2
under 100% of light intensity

0 without hydrogen peroxide
A with 0.001 M

Percentage of Remaining

 

0 10 20 30 40 50 60

Time (min)
Figure A.22 Photodegradation of 2-chlorophenol in UV/TiOz/Air/NaHCO3

under 100% of light intensity

99

Tabla A.23 Data af Figpra Ag fram axmrimantal systam No. 15 and Na. 16

 

UV/TiOZ/Air/NaHCO3 with and without l-IZO2
under 70% of light intensity

0 without hydrogen peroxide
100 A with 0.001 M

80
60

40

Percentage of Remaining

20

 

0 10 20 30 40 50 60

Time (min)
Figure A.23 Photodegradation of 2-chlorophenol in UV/TiOz/Air/NaHCO3

under 70% of light intensity

100

Tabla A.24 Data of Figara A.24 fram axparimantal sysgm No. 15 Ed No. 16

 

UV/TiOZIAir/NaHCO3 with and without 11302
under 50% of light intensity

0 without hydrogen peroxide

1 00 A with 0.001 M hydrogen peroxide

80
60

40

Percentage of Remaining

20

 

0 10 20 30 40 50 60

Time (min)
Figure A.24 Photodegradation of 2-chlorophenol in UV / Ti02 / Air/ NaHCO3

under 58% of light intensity

101

Table A.25 Data af Eigara A25 fram axparimantal systam No. 15 and No. 16

 

UV/Ti02/Air/NaHCO3 with and without F502
under 35% of light intensity

0 wiyhout hydrogen peroxide
A with 0.001 M hydrogen peroxide

80
60

40

Percentage of Remaining

20

 

0
0 1O 20 30 40 50 60

Time (min)
Figure A.25 Photodegradation of 2—chlorophenol in UV/TiOz/Air/NaHCO3
under 35% of light intensity

Ta le A.2

Experimental Illuminated .
s stems Method

nst f 2

102

101' hnl

lv

om uter

Rate constants

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

a: Concentration of hydrogen peroxide in this Table is 0.001 M.

 

 

 

(1 / min)
UV /Ti02 Continuous None 0.1 15
UV/TiOz/ Air Continuous None 0.125
UV / Ti02 / 02 Continuous None 0.145
UV /Ti02 / H2023 Continuous None 0.235
UV/TiOz/Oz Intermittent 7.58 0.056
UV / TiOz/ H202 Intermittent 719 0.086
UV /Ti02 Intermittent 3 0.109
5 0.082
6.5 0.067
8 0.064
10 0.026
UV [1302 /NaHC03b Continuous 7.51 0.061
UV/TiOz/Air/NaHC—O3 Continuous 3.6 0.047
5 0.066
6.5 0.070
8 0.061
10 0.030
UV/TiOz/Air/NaHCO3/H202'q—I'Continuous 3.6 0.106
5 0.108
6.5 0.085
8 0.078
10 0.036
UV/Ti02/Air/NaHC03/Fei+ Continuous 6.87 0.044
UV / Ti02 / Air / NaHCO3 / Fe2+ Continuous 6.93 0.064
‘ /H202
UV / 1102/ Air/ NaHC—Os mtinuous
Light intensity
1 00% 6.74 0.078
70% 6.84 0.062
58% 6.86 0.056
35% 6.92 0.046
UV/ Ti02 / Air/ N aHCO3 / H202 Continuous
Light intensity
1 00% 6.85 0.142
70% 6.87 0.093
58% 6.86 0.077
35% 6.96 0.048