Removal of Lindane from Water Using Advanced Oxidation Techniques

Home , Polydimethylsiloxane, Saleem Khan

REMOVAL OF LINDANE FROM WATER USING ADVANCED OXIDATION TECHNIQUES

BY SANAULLAH KHAN

NATIONAL CENTRE OF EXCELLENCE IN PHYSICAL CHEMISTRY, UNIVERSITY OF PESHAWAR−PAKISTAN JULY, 2014

REMOVAL OF LINDANE FROM WATER USING ADVANCED OXIDATION TECHNIQUES

A dissertation submitted to the University of Peshawar in partial fulfillment of the requirement for the degree of

DOCTOR OF PHILOSOPHY IN PHYSICAL CHEMISTRY

BY SANAULLAH KHAN

NATIONAL CENTRE OF EXCELLENCE IN PHYSICAL CHEMISTRY, UNIVERSITY OF PESHAWAR, PESHAWAR, PAKISTAN JULY, 2014

NATIONAL CENTRE OF EXCELLENCE IN PHYSICAL CHEMISTRY, UNIVERSITY OF PESHAWAR, PESHAWAR, PAKISTAN JULY, 2014 It is recommended that the dissertation prepared by SANAULLAH KHAN entitled “Removal of Lindane from Water using Advanced Oxidation Techniques” be accepted as a fulfillment of the requirement for the degree of DOCTOR OF PHILOSOPHY IN PHYSICAL CHEMISTRY APPROVED BY

(PROF. DR. HASAN M. KHAN) (PROF. DR. M. SALEEM KHAN)

Research Supervisor Director

______

External Examiner Internal Examiner

DEDICATION

My thesis is dedicated to the most beloved personality of Allah (Subhana-hu Wa

Ta’ala), Prophet Mohammed (PBUH). I dedicate my effort of writing my dissertation to my respectable father whose continue love and kindness, and encouragement always helped me to keep my determination steadfast.

Sanaullah Khan

TABLE OF CONTENTS

Title Page No Title i Table of Contents vi Acknowledgements x Abstract xii Keywords xiv List of abbreviations xv List of publications xviii 1. INTRODUCTION 1 1.1 Water Pollution Issue 1 1.2 Persistent Organic Pollutants (POPs) 4 1.3 Lindane Contamination 6 1.3.1 Uses of lindane 8 1.3.2 Toxicity of lindane 9 1.4 Literature Review 11 1.4.1 Lindane degradation by various methods 11 1.4.1.1 Physical methods 11 1.4.1.2 Biological methods 13 1.4.1.3 Advanced oxidation processes (AOPs) 14 1.4.1.3.1 Fenton and photo-Fenton reactions 15

1.4.1.3.2 UV/H2O2 process 16 1.4.1.3.3 Ozonation 17 1.4.1.3.4 Reduction by zero-valent iron 19 1.4.1.3.5 Electrolysis 19 1.4.1.3.6 Microwave irradiations 20 1.4.1.3.7 Ionizing radiations 20

1.4.1.3.8 TiO2 photocatalysis 22 1.4.1.3.9 Sulfur radical-based AOPs 24 1.5 Aims and objectives of the Present Work 25 6

2. EXPERIMENTAL 27 2.1 Materials 27 2.2 Sample Collection 28 2.3 Preparation of Solutions 30 2.4 Extraction Technique 30 2.5 Reactor Design 31 2.5.1 Gamma rays reactor 31 2.5.2 UV reactor 32 2.5.3 UV/Vis reactor 32 2.5.4 Solar reactor 34 2.5.5 Fluorescence light reactor 34 2.5.6 Dark reactor 35 2.6 Calibration of Radiation Sources 35 2.6.1 Calibration of gamma radiation source 35 2.6.2 Calibration of UV radiation source 38 2.7 Qualitative/Quantitative Analysis of Lindane and other Products 40 2.7.1 Gas chromatography-electron capture detector (GC/ECD) 40 2.7.2 Gas chromatography-mass spectrometry (GC/MS) 40 2.7.3 Ion chromatography (IC) 41 2.7.4 High performance liquid chromatography (HPLC) 42 2.7.5 Total organic carbon (TOC) analyzer 42 2.7.6 UV spectrophotometer 42

2.8 Synthesis of Sulfur Doped TiO2 Photocatalyst 43

3. RESULTS AND DISCUSION 49 3.1 Determination of Lindane in Field Water Samples of District Swabi, KP, Pakistan 49 3.1.1 Optimization of the GC/ECD method for lindane analysis 49 3.1.2 Precision, accuracy, reproducibility and relative recoveries of the method 49 3.2 Application of the optimized GC/ECD method to Real Water Samples 54

3.3 Gamma Radiation-induced Degradation of Lindane in Water 57 3.3.1 Kinetic Studies of the Gamma Radiation-induced Degradation of Lindane 58 3.3.1.1 Effect of initial solute concentration 59 3.3.1.2 Effect of solution pH 62 3.3.2 Scavenging effects on gamma radiation-induced degradation of aqueous lindane 65

3.3.2.1 Radiolysis of lindane in N2-saturated solution (control experiments) 66 3.3.2.2 Radiolysis of lindane in aerated solution (natural condition) 66

3.3.2.3 Radiolysis of lindane in N2-saturated 60 mM i-PrOH solution (reductive conditions) 66

3.3.2.4 Radiolysis of lindane in N2O-saturated solution (oxidative conditions) 67 3.3.3 Role of individual reactive species in lindane degradation 69 3.3.4 Dechlorination studies of lindane 75 3.3.5 Effect of common inorganic ions on lindane degradation 76

3.3.6 Effect of H2O2 on lindane degradation 79 3.3.7 Effect of natural organic matter (NOM) and synthetic organic pollutants 79 3.3.8 Pulse radiolysis of lindane–Hydrated electron rate constants − (eaq + lindane) 82 3.3.9 Competition kinetics–Hydroxyl radical rate constant (•OH + Lindane) 83 3.3.10 Variation of solution pH during irradiation 84 3.3.11 Identification of by-products and possible reaction pathways 86 3.3.12 Removal efficiency of lindane 89 3.4 Degradation of Lindane by Photochemical Oxidation 100 3.4.1 Degradation of lindane by solely PMS or ferrous iron (Fe2+) and direct UV photolysis 100 3.4.2 PMS activated by ferrous ion: Fe2+/PMS system 103

3.4.3 Fe2+/PMS system assisted by tube-light radiation: Tube-light/Fe2+/PMS system 105 3.4.4 PMS activated by UV radiation: UV/PMS system 108 3.4.5 UV/Fe2+/PMS system 113 3.4.6 Kinetics of UV/PMS oxidation 117 3.4.6.1 Effect of initial concentrations of lindane 118 3.4.6.2 Effect of solution pH 122 3.4.6.3 Effect of humic acid on the UV/PMS process 123 3.4.6.4 Effect of inorganic anions on the UV/PMS process 126 3.4.7 Mineralization study 129 3.4.8 Oxidant residue analysis of UV/Fe2+/PMS process 135 3.4.9 Product analysis and reaction mechanism of the UV/PMS system 138 3.4.10 Degradation of trichlorobenzene (TCB) by photochemical oxidation 144

3.5 Hydroxyl radical based AOP using UV activated H2O2 for lindane degradation 150

3.5.1 Comparison of the UV/PMS system with UV/H2O2 system 152

3.6 Photocatalytic activity of sulfur doped TiO2 (S-TiO2) and TiO2 under Visible light 157

3.6.1 Factors affecting efficiency of photocatalytic activity of S-TiO2 Process 161 3.6.1.1 Effect of solution pH 161 3.6.1.2 Effect of initial concentration of lindane 162 3.6.1.3 Effects of inorganic oxides. i.e., PS and PMS on the degradation

of lindane in the simulated solar/S-TiO2 system 168 4. CONCLUSIONS AND FUTURE PERSPECTIVES 172 4.1 Conclusions 172 4.2 Future Perspectives 174 REFERENCES 176

ACKNOWLEDGEMENTS

My first and foremost gratitude goes to ALLAH (Subhanahu Wa Ta’ala)”, who enabled me to finish this great job, which otherwise was very difficult for me to do without His Wills. My beloved prophet “Mohammad (Peace be upon him)”, whose deep love always inspired me, deserves highest gratitude after ALLAH (SWT). Next, I would like to present my great appreciations to my kind supervisor Prof. Dr. Hasan M. Khan, former Director of the Centre, for his motivation, encouragement, cooperation, and friendly behavior. His critical discussion on many research issues always inspired me and led me to submit my thesis in the present form. I am very thankful to Prof. Dr. M. Saleem Khan, Director, National Center of Excellence in Physical Chemistry, University of Peshawar, for his sympathetic attitude and provision of all sorts of facilities during the course of my Ph.D research work. I am really grateful to Prof. Dr. Dionysios D. Dionysiou, University of Cincinnati, Cincinnati, USA, for his help, guidance, and for providing all kinds of research facilities during six month research work at the University of Cincinnati. I am truly thankful to Xuexiang He and Changseok Han, Ph.D research students in the University of Cincinnati, for their sincere help in the research work during my stay at the University of Cincinnati. I am grateful to my colleagues, Drs. Javed Ali Khan, Murtaza Sayed and Noor Samad Shah for their help and cooperation in many ways, especially in experimental work as well as in results and discussions. My sincere appreciations also go to my other lab fellows, Mr. Shah Nawaz, Dr. M. Ismail, Faiza Rehman, Fazl-e-Hadi, Fayaz Ali, Razaullah, and Jehangir Khan for their friendship and necessary help. I really enjoyed their company during the course of my Ph.D study. My good wishes go to all staff members of the National Center of Excellence in Physical Chemistry, University of Peshawar for their respective help and support. I am also thankful to the staff members of the Nuclear Institute for Food and Agriculture (NIFA), Peshawar, specially the Director, Dr. Ihsanullah and Mr. Tariq Nawaz for providing research facilities at NIFA, particularly the gamma irradiator.

Finally, my cordial thanks and appreciations go to my parents, brothers especially Hamidullah Khan and Shafiullah Khan, uncle, Prof. Ghani-ur-Rehman, sisters, wife and my friends particularly Zia-ur-Rehman, Tariq Mahmood and Shams-ul-Haq, for their infinite love, sympathy, motivation, encouragement and prayers.

The Higher Education Commission of Pakistan (HEC) is highly acknowledged for providing financial assistance through 5000 Indigenous Ph.D fellowships throughout the Ph.D degree program. HEC is also acknowledged for granting fellowship under the International Research Support Initiative Program (IRSIP), to carry out a part of my Ph.D research work at the University of Cincinnati, OH, USA. Sanaullah Khan

ABSTRACT

Organochlorine pesticides are highly persistent and the most emerging endocrine disrupting chemicals in the environment. In this study, several emerging advanced oxidation processes (AOPs), i.e., gamma ionizing radiation, sulfate radical based-AOPs, and non-metal doped TiO2 photocatalysis were investigated for the degradation of lindane in aqueous solution. The effects of water quality and process parameters, such as solution pH, initial concentration of pollutant, initial concentration of oxidant and/or catalyst, oxidant/catalyst molar ratios, presence of inorganic ions and natural organic matters, were studied. All of the studied AOPs showed high efficiency, and ultimately led to complete degradation of lindane under different conditions. The degradation of lindane by the studied AOPs followed pseudo-first order kinetics. The observed pseudo-first order rate constant decreased while degradation rate increased with the increasing initial concentration of lindane. Different initial pH presented a different effect on lindane degradation. The highest efficiency of lindane degradation was achieved at neutral pH.

The presence of co-existing background constituents, such as natural organic matter

2− − 2− − − − and/or inorganic anions (CO3 , HCO3 , SO4 , Cl , NO3 , and NO2 ) presented a different effect on lindane degradation, based on the competition of these constituents with lindane for the reactive radicals generated in various AOPs. The results were

• •− discussed in terms of reactivity of hydroxyl radical ( OH), sulfate radical (SO4 ) and

− hydrated electron (eaq ) with lindane. Hydrated electron was found to be the most reactive species, as suggested from its high second-order rate constant of 1.26 × 1010 M˗1 s−1 (as

9 8 −1 −1 •− • compared to 1.3 × 10 and 6.8 × 10 M s for SO4 and OH, respectively), determined via competition kinetics method. Based on the detected by-products through GC-MS

analysis, a plausible reaction mechanism was proposed for UV activated

− peroxymonosulfate process (i.e., UV/HSO5 ), suggesting dechlorination, chlorination,

• •− dehydrogenation and hydroxylation via OH and/or SO4 attack. The studied AOPs also showed significant efficiency in dechlorination and mineralization of lindane. The

− addition of 0.2 mM peroxymonosulfate (HSO5 ) showed a large enhancing effect on visible and solar light-assisted sulfur-doped titanium dioxide (S-TiO2) photocatalysis of

• •− lindane. The reaction mechanism revealed that OH and SO4 reacted with lindane via

− hydrogen abstraction pathway, while eaq followed dissociative electron capture mechanism. Trichlorobenzene (TCB), a typical reaction by-product of lindane, was readily degraded by various UV/oxidant processes tested. Assessment of lindane residues in the surface water samples, in different regions of district Swabi, Khyber

Pakhtoonkhwa (Pakistan), was carried out; indicating thirteen out of the eighteen samples analyzed, were contaminated by varying amounts of lindane. The outcome of this study will provide useful scientific information on the effectiveness of gamma radiations, S-

TiO2 photocatalysis, and various sulfate radical based-AOPs on the degradation of organic compounds, especially problematic organochlorine pesticides.

Keywords: Lindane; Gamma radiations, Advance Oxidation Processes (AOPs);

UV/Peroxymonosulfate; S-TiO2 photocatalysis; Rate constants; Water treatment.

LIST OF ABBREVIATIONS

AOPs Advance oxidation processes AOTs Advanced oxidation technologies ATSDR Agency for toxic substances and disease registry a.u arbitrary unit BET Brunauer-Emmett-Teller BOD Biological oxygen demon CNS Central nervous system COD Chemical oxygen demon CP Chlorophenol CSTR Continuous stirred tank reactor DBCP 1, 2-dibromo-3-chloropropane DBPs Disinfection by-products EDCs Endochrine disrupting chemicals EI Electron impact FBR Fix bed reactor FA Fulvic acid FSD Food and Soil Division GAC Granular activated carbon GC/ECD Gas chromatograph electron capture detector GC/MS Gas chromatograph mass spectrometry HA Humic acid HCB Hexachlorobenzene HCH Hexachlorocyclohexane HPLC High performance liquid chromatograph IC Ion chromatography IERC International agency for research on cancer KP Khyber Pakhtoonkhwa LOD Limit of detection

LOQ Limit of quantification MAL Maximum acceptable level MC Mechanochemical MOAFF Ministry of Agriculture, Forestry, and Fisheries MTBE Methyl tert-butyl ether MW Microwave ND Not detected NIFA Nuclear institute for food and agriculture NIST National institute of standards and technology NOMs Natural organic matters OCPs Organochlorine pesticides PAC Powdered activated carbon PCBs Polychlorinated biphenyls PCE Tetrachloroethene PCPs Pharmaceuticals and personal care products PF Photo-Fenton PMDS Polydimethylsiloxane PMS Peroxymonosulfate POPs Persistent organic pollutants PS Persulfate RO Reverse osmosis ROSs Reactive oxygen species RSD Relative standard deviation SPME Solid phase micro extraction SLA Solar light activity SR/AOPs Sulfate radical-based advanced oxidation processes S/N signal-to-noise ratio TCE tetrachloroethylene TOC Total organic carbon TTIP Titanium (IV) isopropoxide

UNEP United nation environmental programme UNO United nation organization US EPA United state environmental protection agency UV Ultra Violet VLA Visible light activity WHO World Health Organization ZVI Zero-valent iron

LIST OF PUBLICATIONS

1. S. Khan, X. He, H.M. Khan, D. Boccelli, D.D. Dionysiou, Efficient degradation

of lindane in aqueous solution by iron (II) and/or UV activated

peroxymonosulfate, Journal of Photochemistry and Photobiology A: Chemistry,

316 (2016) 37-43.

2. H.M. Khan, A.A. Khan, S. Khan, Application of DNA comet assay for detection

of radiation treatment of grams and pulses, 48 718-723.

3. M. Sayed, M. Ismail, S. Khan, S. Tabassum, H.M. Khan, Degradation of

ciprofloxacin in water by advanced oxidation process: kinetics study, influencing

parameters and degradation pathways, Environmental technology, (2015) 1-13.

4. M. Sayed, M. Ismail, S. Khan, H.M. Khan, A comparative study for the

quantitative determination of paracetamol in tablets using UV-Visible

spectrophotometry and high performance liquid chromatography, Physical

Chemistry 17(1) (2015) 1-5.

5. S. Khan, C. Han, H.M. Khan, D.L. Boccelli, D.D. Dionysiou, Sulfur-doped TiO2

for photocatalytic degradation of lindane under visible and solar light irradiation:

Strong enhancement due to peroxymonosulfate addition (submitted to Chemical

engineering journal).

6. S. Khan, X. He, J.A. Khan, H.M. Khan, D.L. Boccelli, D.D. Dionysiou, Kinetics

and mechanism of sulfate radical- and hydroxyl radical-induced degradation of

highly chlorinated pesticide lindane in UV/peroxymonosulfate system (submitted

to Chemical engineering journal).

1. INTRODUCTION

1.1 Water Pollution Issue

Water is a critical resource and basic necessity for each and every individual on land.

The shortage of freshwater resources worldwide will only become more critical as the world population increases and the climate changes. The rapid progress in the field of agriculture and industrialization though contributed a lot to the economic development of nations; it has resulted in heavy losses to society in terms of water, soil and air pollution

(Reddy and Behera, 2006). Groundwater and surface waters that constitute an integral part of the continental water cycle can transport and spread contaminants from spatially limited industrial areas or mining areas to extensive downstream regions. This is in addition to spreading of more diffuse contamination, such as from pesticides distributed over agricultural fields and vehicular transport system. Water resources are affected by man-made pollution worldwide to such an extent that restoration to pristine conditions is not achievable (Tornqvist et al., 2011).

The significance of accessibility of good quality drinking water can be recognized by the press release of UNO Secretary General on world water day 2002: “An estimated

1.1 billion people lack access to safe drinking water, 2.5 billion people has no access to proper sanitation, and more than 5 million people die each year from water-related diseases–10 times the number killed in wars, on average, each year. All too often, water is treated as an infinite free good. Yet even where supplies are sufficient or plentiful, they are increasingly at risk from pollution and rising demand. By 2025, two thirds of the world's population is likely to live in countries with moderate or severe water shortages”

(UNO, 2002). Water is the most essential element for existence of all kinds of life on the

Earth. Most part of the world’s water resources is confined to occasions as saline water; only 3% is available as freshwater. Out of the total amount of freshwater, only 0.01% is available for human use (Hinrichsen and Tacio, 1997). Unfortunately, this small proportion of freshwater is constantly polluted due to rapid population growth, urbanization and industrialization and gross agriculture activities. According to UNO report, population of the world is regularly increasing, while the existing fresh water resources are declining day by day. As a result, many nations of the world, particularly some countries in Middle East, Africa, and South Asia will have severe problems of water shortage in the next two decades (UNO, 2002).

Although water contamination is a global issue, developing countries are facing extra problems because of deprived management and poor monitoring policy. Like other developing nations, there is severe public health concern in Pakistan due to low quality water. Pakistan ranks at number 80 among 122 nations regarding drinking water quality.

Drinking water sources, both surface and groundwater are contaminated with bacteria, toxic metals and pesticides throughout the country. Various drinking water quality parameters set by the World Health Organization (WHO) are frequently violated.

Different anthropogenic sources, like improper disposal of municipal and industrial effluents and indiscriminate usage of pesticides for agriculture purposes are primarily involved in the deterioration of water quality. Microbial and chemical pollutants are the main factors responsible exclusively or in combination for various public health problems

(Azizullah et al., 2011). The country has essentially exhausted its available water resources; it is considered as water stressed and is likely to have a water scarcity in the near future (Hashmi et al). The water precipitation rate is lower than the evaporation rate

in the country. This causes a continuous decrease in water quantity in its rivers, lakes and diminishing the groundwater as well. The problem is further provoked by factors like long droughts and lack of construction of new water reservoirs (Ullah et al., 2009). This decrease in water quantity coupled with increasing demand resulted in severe water shortage in almost all sectors of the country. Large number of organic recalcitrants, including solvents, pesticides, paints, dyes, phenols, petrochemicals, polymers, pharmaceuticals and personal care products (PCPs) are continuously added into the water resources since the early development in the field of industrialization. All these chemicals are highly toxic to animals, and many of these substances can be readily absorbed through the skin. These factors can badly affect higher classes of mammals including human beings (Sheoran, 2008a).

Chemical pollution of surface waters presents a hazard to the aquatic environment as well as a threat to human health. Among various chemicals, pesticide contamination of surface water resources is a major water quality issue throughout the world.

Pesticide is a general term for substances, which are used to poison pests (weeds, insects, molds, rodents etc.). The pesticides most acutely dangerous to humans are insecticides and rodenticides. Synthetic pesticides have been popular with farmers, because of their efficacy and cost effectiveness, but they also have huge environmental costs (Agrawal, 2010). Due to new arrivals in farming practices and intensive revolution in the field of agriculture, the worldwide production and consumption of pesticide has significantly increased in recent years. The extensive use of pesticides for agricultural and several other purposes has resulted in the occurrence of pesticides residues in various environmental segments of the world. Pesticide residues reach the aquatic environment

through direct run-off, leaching, careless disposal of empty containers, equipment washing, etc. Pesticide contamination of surface waters has been well documented worldwide and constitutes a major issue that gives rise to concerns at local, regional, national and global scales (Planas et al., 1997; Huber et al., 2000). So, it is clear that the strategy to find solutions to this immense problem by developing cheap and green remediation technologies is very important.

1.2 Persistent Organic Pollutants (POPs)

Persistent organic pollutants (POPs) are carbon-based synthetic organic products and by-products of priority concern for community of environmental scientists and engineers, worldwide (Tanabe et al., 1994; Hansen 1998). POPs constitute a diverse group of organic substances with some intrinsic physical-chemical properties, such as toxicity, persistency, bioaccumulative and long-range transportable nature, which dictate their environmental behavior (Wania, 2003, 2005). There are many thousands of persistent chemicals, often coming from certain series or `families' of chemicals (e.g. there are theoretically 209 different polychlorinated biphenyls, differing from each other by level of chlorination and substitution position). There is no fixed opinion about the half-life of a substance in a particular media to be conferred `persistent'; however, generally a POP has a half-life of several days in the atmosphere and several years in the soil and sediment (Jones and de Voogt, 1999). POPs are typically `water-hating' and `fat- loving' chemicals. So, they prefer lipids rather than the aqueous environment of cells inside living organisms and get stored in fatty tissue. POPs have the tendency to enter the gas phase under environmental temperatures. Hence, they may volatilize from water

bodies, soils and vegetation into the atmosphere. Low POP levels might be increased by biomagnifications through the transmission process in the food chain (Tieyu et al., 2005).

Environmental occurrence of POPs is a global rather than a regional problem, because the POPs used in tropical regions will be carried by long-range atmospheric transport to polar and environmentally pristine areas, such as the Arctic region (Hargrave et al., 1988; Patton et al., 1989; Barrie et al., 1992). As a result, POPs pollution has touched every region in the world. POPs have been reported to cause variety of effects including immunologic, teratogenic, carcinogenic, reproductive, developmental, behavioral, neurological and endocrine problems in organisms (Pages et al., 2002).

Because of massive adverse effect on environment, a great deal of attention is paid to

POPs contamination problems, and strong action has been taken by many developed and developing nations (Tieyu et al., 2005). In 2004, the Stockholm Convention on POPs

(POPs Treaty) that was adopted by the United Nations Environmental Programme

(UNEP) came into force and aims to protect human health and the environment from negative effects of POPs by reducing or eliminating releases of POPs to the environment.

At first, the treaty promotes the global regulations on the production and usage of twelve substances, the so-called “dirty dozen” which consists of eight kinds of pesticides, including dieldrin, aldrin, endrin, chlordane, heptachlor, DDT, toxaphene, mirex, two kinds of industrial chemicals (polychlorinated biphenyls (PCBs) and hexachlorobenzene

(HCB)), and two kinds of by-products (polychlorinated dibenzofurans and polychlorinated dibenzo-p-dioxins) (Hosomi, 2001, 2002, UNEP, 2001). In 2009, hexachlorocyclohexane (HCH) among eight other compounds, was added as new

Stockholm Convention POPs, thus extending the original POPs list (12) to a total of 21 members (Vijgen et al., 2011).

Organochlorine pesticides (OCPs) are one of the most important persistent organic pollutants (POPs) and have been of great concern around the world owing to their chronic toxicity, persistence and bioaccumulation (Willett et al., 1998). OCPs are ubiquitous anthropogenic environmental contaminants posing great threats to ecosystems and human health (Kalajzic et al., 1998). Although the application of these chemicals has been banned or restricted in many countries, especially the developed ones, some developing countries are still using these compounds because of their low cost and versatility in industry, agriculture and public health (Tanabe et al., 1994, Nasir et al., 1998). Among the OCPs, lindane has been shown to be very persistent, bioaccumulative and toxic to humans (Moore and Walker., 1964, Sarkar et al., 2008).

1.3 Lindane Contamination

Lindane is an organochlorine insecticide which has been extensively used on wide varieties of fruit, grain and vegetable crops and conifer trees for control of leaf-eating insects and pests (Ware and Whitacre, 2004). Lindane is also known as gamma- hexachlorocyclohexane (γ-HCH) since it is made up of at least 99% of the gamma isomer of hexachlorocyclohexane (HCH). The chemical structure of lindane is shown in Figure

1.1. HCH is available in two formulations: technical HCH and technical lindane. A total of eight HCH isomers denoted by Greek letters α, β, γ, δ, ε, ζ, η and θ have been identified in technical HCH. However, only the α, β, γ, δ, and ε-isomers, varying in the following percentages: α: 60-70%, β: 5-12%, γ: 8-15%, δ: 6-10%, and ε: 3-4%, are stable isomers of technical HCH, while lindane is almost pure γ-isomer (Rankenberger., 2002).

Hexachlorocyclohexane was initially synthesised in 1825, but the pesticidal properties of

HCH were recognized in 1942 (Metcalf., 1955). Commercial production of technical- grade HCH began in 1943 by Imperial Chemical Industries Ltd; UK (ICI) and use started up in the United Kingdom (Li et al., 2011). HCH can be synthesised by photochlorination of benzene. When the resulted product is treated with acetic acid or methyl alcohol,

99.9% pure γ-HCH isomer (i.e., lindane) is obtained upon fractional crystallization process. The γ-isomer (CAS Registry No. 58-89-9) is the isomer with the highest pesticidal activity; however, technical mixtures of all isomers (CAS Registry No. 608-73-

1) have been widely used as commercial pesticides (Metcalf., 1955). Total global production and use of the different HCH isomers is difficult to determine, and estimates vary considerably. Voldner and Li estimated total use of technical-grade HCH and lindane to be 550 000 and 720 000 metric tons, respectively (Voldner and Li., 1995).

However, later calculations by Li and co-workers placed total cumulative world consumption of technical-grade HCH as high as 6 million metric tons (Li et al., 1998).

More recently, total global usage of technical HCH between 1948 and 1997 was estimated to be around 10 million metric tons (Li., 1999). Environmental and human health concerns led to the banning of technical-grade HCH in many countries during the

1970s. China, India, and the former Soviet Union remained the largest producers and users of HCH in the early 1980s. In 1980, the annual consumption of technical HCH in two Asian countries, India and China, accounted for more than 84% of the total technical

HCH consumption in the world. Of the 90 000 tons applied in 1990, 51 000 tons was used in India. In China, the production of HCH was banned in 1983, but residual stocks may have been used until 1985 (Li et al., 1998). In1990, although increased in India, the

usage of technical HCH decreased dramatically among other countries. In 1990, production of technical-grade HCH was also prohibited in the former Soviet Union, and use of residual stocks was restricted to public health and specific crop uses. Literature survey shows that the leading lindane consuming countries include both developing countries as well as developed countries. In 1990, France, Italy, Niger, Canada, the

United States, India and China were among the leading lindane consuming countries with annual usage more than 100 tons (Li et al., 1996).

1.3.1 Uses of lindane

Lindane has been released to the environment during its formulation process and through its use. Lindane is used as an insecticide on fruit, grain and vegetable crops, in warehouses and in public health measures to control insect-borne diseases (Li et al.,

1998). Together with fungicides, lindane can also be used as a seed treatment agent for barley, corn, oats, rye, sorghum, and wheat. Lindane is also used in a variety of domestic and agricultural applications, such as dips, sprays and dust for livestock and domestic pets (Li., 1999). The forestry industry also uses lindane to control pests on cut logs

(Donald et al., 1997). Lindane is also used topically for the treatment of head and body lice and scabies; it is available in 1 percent preparations as a lotion, cream, or shampoo

(Nordt and Chew., 2000). The various formulations of lindane and technical HCH have many trade names, including Agrocide, Ben-Hex, Gammexane, Kwell, Quellada,

Lindatox, and Tri-6 (Sang et al., 1999). Extremely low cost of lindane led to its wide use, particularly in some developing countries (Abhilash and Singh, 2009). Lindane residues persist in the environment, undergo volatilization in tropical conditions, migrate to long distances with air current, deposit in colder regions, and cause widespread contamination;

the γ-HCH-residues reach human body via food chain and get biomagnified at each trophy level (Benimeli, 2008). Lindane contamination has been detected in air, surface water, groundwater, sediment, soil, fish and other aquatic organisms, wildlife, food, and humans (Iwata et al., 1994, Li, 1999). Human exposure results primarily from medicinal use and from ingestion of contaminated plants, animals, and animal products. Lindane has not been found to be a major contaminant of drinking water supplies in U.S.A

(ATSDR, 2005)

The commercial application of lindane was launched in the decade following

World War II and was used extensively in Europe, U.S.A., and other developed countries in 1970s. Lindane is considered to be highly toxic to aquatic organisms, and moderately toxic to birds and mammals. People who are occupationally exposed to it are advised to avoid its contact with eyes, skin and via inhalation. Symptoms of acute toxicity in humans include mild skin irritation, headache, dizziness, seizures, diarrhea, nausea, vomiting and even convulsions as well as effects on the gastrointestinal tract, cardiovascular and musculoskeletal systems (Sang et al., 1999).

1.3.2 Toxicity of lindane

Lindane is a well known reproductive and developmental toxicant with a pattern of effects suggesting an endocrine-related mechanism (Pages et al., 2002, Traina et al.,

2003). Lindane can inhibit the synthesis of DNA, RNA, protein and carbohydrate contents in animals (Al-Chalabi and Al-Khayat, 1989). It has adverse effects on the central nervous system (CNS) and has teratogenic, immunotoxic, and neurotoxic properties (Hayes, 1982). Acute and chronic exposure of lindane has been shown to produce marked neurological and hepatotoxic effects in experimental animals (Oesch,

1982). It produces neurotoxic, hepatotoxic and uterotoxic effects in rats (Criswell and

Loch-Caruso, 1999). The neurotoxicity of lindane involving different manifestations of hyperexcitability and, at high doses, convulsions have been extensively reported (Wooley et al., 1985, Tusell et al., 1988).

With this comment, lindane is considered as a possible human carcinogen by the

United States Environmental Protection Agency (US EPA) and the International Agency for Research on Cancer (PHG, 1999). WHO has classified lindane as a “moderately hazardous” substance (WHO, 2004).

Figure 1.1. Chemical structure of lindane

1.4 Literature Review

Lindane and other HCHs have long been a well studied class of organochlorine compounds with respect to environmental fate and effects. Lindane has been used extensively in both agricultural and pharmaceutical commercial applications for more than fifty years (Walker et al., 1999, Sarkar et al., 2008). Residues of lindane have been widely detected in environmental water, sediment, and organism samples. Because of potential water contamination, there is a need to develop new technological approaches for rapidly destroying lindane and other POPs in aqueous solution.

1.4.1 Lindane degradation by various methods

Despite its persistence, lindane degradation has been reported using various conventional methods as well as advanced oxidation technologies (AOTs). The conventional decontamination techniques consist of physical methods (adsorption on activated carbon, coagulation-flocculation, mechanochemical treatment, membrane technology and reverse osmosis) and biological methods i.e. biotransformation.

1.4.1.1 Physical methods

Adsorption process involves the separation of a substance from one phase accompanied by its accumulation at the surface of other. Activated carbon is a highly porous, amorphous solid with large micropore volume and high surface area and adsorption onto activated carbon is considered an advanced water treatment process for the removal of aqueous-dissolved organic pesticides. Both activated carbon types, i.e. granular activated carbon (GAC) and powdered activated carbon (PAC) can be used for the adsorption of pesticides. However, GAC is mainly used due to the advantages of relatively lower carbon application rates, easier handling and the possibility for 29

regeneration of adsorbent (El-Dib, et al., 1975, Pirbazari, et al., 1991). Thacker and co- workers studied the adsorption of aqueous lindane on granular activated carbon (GAC), using two different types of reactors i.e. a continuous stirred tank reactor (CSTR) and fix bed reactor (FBR) systems. The carbon dose required to treat raw water having initial concentrations of 5-10 μg/L of lindane to <2 μg/L was 0.8 and 0.9 g/L for (CSTR) system and 0.5 and 0.6 g/L for (FBR) systems, respectively. This study provides useful results for reducing the concentration of lindane down to the potable level (<2 μg/L), but it does not give information about the dynamics of the system (Thacker et al., 1997). Kouras and co-workers studied the adsorption of lindane from aqueous solutions onto powdered activated carbon (PAC) in batch experiments. The results indicated that PAC doses greater than 20 mg/L were found to be necessary in order to reduce lindane from initial concentration of 10 mg/L down to 0.1 mg/L within 1 h contact time. pH of the solution did not show any effects on the adsorption efficiency (Kouras et al., 1998).

Membrane technology comprising of microfiltration or ultra filtration process has been remarked among innovative technologies, because no chemical agents used and high quality water is constantly produced with simple automation of process (Sadr

Ghayeni et al., 1996, Comerton et al 2005, Wintgens et al., 2005).

Reverse osmosis has been gradually finding application in the treatment of a variety of domestic, industrial, and hospital wastewaters (Plakas and Karabelas, 2012).

Chian and co-workers studied the removal of pesticide with reverse osmosis using different kinds of membranes. The efficiency of the process was found to vary with the kind of the membrane used. It was seen that non-cellulosic membranes showed better efficiency for removal of pesticides than cellulosic type membrane (Chian et al., 1975).

Mechanochemical (MC) treatment is another developed technology of pollutant removal approved by the Ministry of Agriculture, Forestry, and Fisheries of Japan

(MOAFF). This technology has successfully been used for degradation various kinds of hazardous chemicals, particularly organochlorine pesticides in recent decades (Hall et al.,

1996, Loiselle et al., 1996, Tanaka et al., 2003, Birke et al., 2004). Nomura and co- workers studied the applicability of the mechanochemical (MC) treatment technology for the destruction of lindane in batch experiments (Nomura et al., 2012). The degradation of lindane ultimately led to the formation of trichlorobenzenes (TCB) as final by-product.

Further degradation products (dichlorobenzenes, monochlorobenzene, and benzene) were also detected but in limited amount. Traces of methane and ethane were also detected, suggesting cleavage of the C–C bonds of the cyclohexane ring and hydrogenation.

1.4.1.2 Biological methods

Bioremediation, which incorporate the useful utilization of microorganism for the degradation of target pollutants, is a competent technique for the biological treatment of industrial wastes and contaminated soils (Alexander, 1994, Crawford and Crawford,

2005). Lindane and other HCHs are known to be susceptible to attack by anaerobic microorganisms that exist in sewage sludge, river or lake sediments, or even in the soil of flooded agricultural fields (Hill and McCarty, 1967, Jagnow, 1977, Middeldorp, 1996).

The degradation mechanism involves reductive dechlorination and dehydrohalogenation pathways to produce less chlorinated cyclohexanes and chlorophenols, and finally gives various tri-, di-, and monochlorobenzenes, and benzene as end products (Quintero, 2005).

Aerobic biodegradation has also been reported, with the data suggesting that lindane could be cometabolized in the environment and, under certain conditions, could serve as

the sole source of carbon for (Sahu et al., 1990, Gupta et al., 2000). Biological methods have shown the feasibility of decreasing the BOD, COD and TOC of wastewater to some extent, and probably they are among the most inexpensive water treatment techniques

(Alinsafi et al., 2005). However, biological methods are inadequate for the treatment of high molecular weight biorefractory compounds, since they are not easily degraded by bacteria and they may also inhibit bacterial development (Georgiou et al., 2003).

1.4.1.3 Advanced oxidation processes (AOPs)

Compared to the processes using mere physical transfer, advanced oxidation processes (AOPs) have merged to be promising alternatives, which successfully transform hazardous organic pollutants into harmless end products. The concept of AOPs was established by Glaze (Glaze, 1987), who defined AOPs as “processes involving the generation of highly reactive oxidizing species capable of attacking and degrading organic substances near ambient conditions of pressure and temperature”. Nowadays,

AOPs are considered as high efficiency physical-chemical processes capable of producing deep changes in the chemical structure of the contaminants via the participation of free radicals (Quiroz et al., 2011). The chemical reactions involved are essentially the same as if the pollutants were otherwise slowly oxidized in the environment, but the oxidation rate is billions of times faster in AOPs (Bolton, 2010).

AOPs are generally based on the use of in situ formed wide variety of highly

• •− • • •− − reactive free radicals, such as OH, SO4 , H , OOH, O2 , e aq etc that can effectively decompose almost all refractory organic pollutants without an additional separation step

(Lin et al., 1995, Anipsitakis and Dionysiou, 2004a, 2004b, Rodriguez et al., 2011 ).

Depending upon the source of the radical generation, various classes of AOPs have been

recognized. The most commonly employed AOPs include Fenton and Fenton-like processes (Fenton, 1894, Matta, 2007), photo-Fenton and photo-Fenton-like reactions

(Bauer and Fallmann, 1997; Kang, 2000; Bandala, 2009), ozonation (Peyton and Glaze,

1988; von Gunten, 2003), direct UV photolysis (Kundu et al, 2005; Dantas, 2010),

UV/H2O2 photolysis (Wallington, 1988, Peternel, 2006), TiO2 photocatalysis (Carey et al., 1976, Malato et al., 2002, Nakata and Fujishima, 2012), sonolysis (Bertelli and Selli,

2004; Vassilakis et al., 2004; Uddin and Hayashi, 2009), radiolysis (Getoff 1995; Cooper,

2003), electrolysis (Martinez-Huitle and Ferro 2006; Radjenović 2012) and reduction by zero valent iron (ZVI) (Kim and Carraway 2000). A short description of the AOPs employed for the treatment of water and wastewater organic pollutants is given below:

1.4.1.3.1 Fenton and photo-Fenton reactions

Among the established AOPs, Fenton and photo-Fenton processes are the most widely studied classes of the reactions for the pollutant degradation, because of (i) easy availability of Fenton’s reagents (ii) cost-effectiveness and (iii) high reactivity and non- selectivity of the generated •OH towards most organic pollutants.

Fenton’s reagent comprises a homogeneous catalytic oxidation system consisting

2+ 2+ of a mixture of hydrogen peroxide (H2O2) and iron (Fe ). The ferrous ion (Fe ) initiates and catalyses decomposition of H2O2 resulting in the generation of hydroxyl radicals,

•OH (Fenton, 1894). The general reaction is:

2+ 3+ • − Fe + H2O2 → Fe + OH + OH (1.1)

In photo-Fenton process in addition to the above reaction, the formation of •OH also occurs by reactions (1.2) and (1.3) (Ho 1986; Kochany and Bolton, 1992):

• H2O2 + hν → 2 OH (1.2)

3+ • 2+ + Fe + H2O + hν → OH + Fe + H (1.3)

Nitoi and co-workers studied simultaneously the degradation of lindane and mineralization of organic chlorine in aqueous solution by photo-Fenton process in batch experiments (Nitoi et al., 2013). The degradation rate followed pseudo-first order kinetics with respect to lindane and organic chlorine mineralization. Results of photo-Fenton reactions assured total organic carbon (TOC) removal with 95% efficiency at 2 h irradiation. Optimal experimental conditions for 99% removal of lindane at initial

−6 −3 2+ concentration of 3.47 × 10 M were: pH = 3, [H2O2] = 29.41 × 10 M, [Fe ] = 3.67 ×

−3 − 10 M, and time = 1 h. The value of klindane higher than kCl confirmed that chlorine liberation did not take place simultaneously with the attack of •OH on lindane. Despite many advantages, however, the requirement of low pH environment is the major limitation to Fenton and photo-Fenton’s processes.

1.4.1.3.2 UV/H2O2 system

• Hydrogen peroxide (H2O2) is one of the most efficient sources of OH radical known for long time, however, the rate of •OH generation is very slow and some activation mean is needed in practical applications. The combination of UV/H2O2 is widely applicable to wastewater treatment for destruction of wide range of toxic pollutants (Wallington, 1988, Peternel, 2006).

Nienow and co-workers (2008) studied the oxidation of lindane in terms of lindane degradation and release of chloride ions in the UV/H2O2 system. Results showed that 90% of the lindane was destroyed in about 4 min under these conditions. In addition, within 15 min, all chlorine atoms were converted to chloride ion, indicating that chlorinated organic by-products do not accumulate in the reaction mixture. The presence

of humic acid (HA) and fulvic acid (FA) showed retarding effects on the degradation process which can be considered as major limitation of the process, since these compounds are commonly found in real water samples. Although the presence of HA and

FA does reduce the reaction rate during treatment, the rate constant in the presence of these compounds remains significantly larger compared to direct photolysis or hydrolysis reactions. They also measured the absolute rate constant for oxidation of lindane with

•OH radical. The results showed that lindane rapidly reacted with •OH and the maximum

−3 −1 reaction rate constant (9.7 × 10 s ) were observed at pH 7 and initial H2O2

• concentration of 1 mM. The OH was generated by UV/H2O2 process, tetrachloroethene

(PCE) was chosen as reference compound and rate constant was determined using competitive kinetics method.

Haag and Yao (1992) determined the rate constants of •OH with lindane molecule by relative rate method, using tetrachloroethylene (TCE) and 1, 2-dibromo-3- chloropropane (DBCP) as reference compounds. A variety of techniques, including

Fenton’s reaction and photo Fenton’s processes, were employed for generation of •OH

9 −1 −1 radical. The absolute rate constant, kabs (lindane) was found to be 1.1 x 10 M s

(Fenton reactions) and 5.2 x 108 M−1 s−1 (photo-Fenton reactions) using TCE and DBCP, respectively, as reference compounds (Haag and Yao, 1992).

1.4.1.3.3 Ozonation

Ozone is long been used for oxidation of many kinds of chemical substances including toxic pollutants (Camel and Bermond, 1998, Javier Benitez et al., 2002). The oxidation of organic pollutant by O3 can be achieved through either direct or indirect pathway. In direct pathway, the organic compound is oxidized by O3 itself in acidic

media whereas indirect ozonation involves the degradation of organics through •OH under basic conditions. Ozone in the gas phase and in solution absorbs ultraviolet radiation with a maximum at 254 nm. In water-rich gas phase, the process involves dissociation into an oxygen molecule and an oxygen atom in lD state. The latter may

• react with H2O to produce two OH via reactions (1.4) and (1.5) (Hoigné 1998):

1 O3 + hν → O2 + O ( D) (1.4)

• O + H2O → 2 OH (1.5)

In the aqueous phase, the radicals apparently combine in a solvent cage to yield hydrogen peroxide (H2O2), which may also photolyze or combine with ozone through a complex radical mechanism. Thus by combining UV radiations with O3, the oxidation power of the system for organic pollutant degradation could be significantly enhanced.

Begum and Gautam (2012) studied the chemical oxidation of lindane with ozone and developed a reaction kinetics and mechanism under various experimental conditions.

Optimization of parameters was done and ozone dose of 57 mg min−1 was chosen as optimal for initial lindane concentration of 25.72 μM, while any further increase in ozone dose had a diminishing effect on the removal efficiency. The pH of the solution was found effective in influencing the extent of degradation and lindane removal rate was favoured under alkaline conditions, plausibly due to higher rate of •OH generation.

Kinetics results showed that the degradation rate follows first-order kinetics with respect to lindane concentration. The observed degradation rate constants (kobs) for initial lindane concentrations of 17.5, 25.72 and 35.00 μM were 0.0243, 0.0333 and 0.056 min−1, respectively.

1.4.1.3.4 Reduction by zero valent iron

Zero-valent iron (ZVI) comprises an emerging technology for the destruction of a large variety of toxic organic pollutants (Agrawal and Tratnyek, 1995, Li et al., 2006).

Schlimm and Heitz (1996) reported dechlorination of lindane to benzene by various transition metals, such as zero valent zinc, Fe/Cu, Al/Cu, Zn/Cu, and Mg/Cu systems.

Wang and co-workers (2009) studied the removal of lindane from water by granular zero- valent iron under varying experimental conditions. The rate of the reaction was found to vary with solution pH, temperature and iron dosage, and the reaction followed pseudo- first-order kinetics with respect to concentration of lindane. The higher reaction temperature, higher ZVI dosage and lower pH favoured the degradation reaction kinetics.

At the end of the reaction, lindane was converted to benzene and chlorobenzenes. Singh and co-workers studied the efficiency of nZVI for the remediation of lindane contaminated soil at different pH values. The reaction followed pseudo first order kinetics leading to complete disappearance of lindane (initial concentration 10 μg g-1) in

24 hours at nZVI concentration of 1.6 g L-1. The system showed higher efficiency at lower pH. Benzene was reported as the final degradation product of lindane under the given conditions (Singh et al., 2011).

1.4.1.3.5 Electrolysis

Electrolysis is a versatile, non-selective and efficient technique employed for destruction of organic pollutants in contaminated water (Martinez-Huitle and Ferro,

2006). The electrochemical methods for the destruction of POPs are mild and environment-friendly and many types of reaction systems can be easily designed. The rate of pollutant decomposition and the removal efficiency depend on the kind of

electrode material, electrolyte composition, pH etc (Comninellis 1994). Electrochemical treatment is successfully applied for many kinds of toxic pollutants chlorinated organic compounds, such as nitrophenol, PCPs and disinfection by-products (DBPs), antibiotics and pharmaceutical compounds (Wei et al., 2011; Patel and Suresh, 2008; Radjenović et al., 2012). Most of the investigators concluded that electrochemical reduction of lindane is a six-electron process that produces benzene as the major product. Beland and co- workers reported that lindane undergoes a one-step, six-electron reduction to afford benzene as the final product at cathode (Beland et al., 1977).

1.4.1.3.6 Microwave irradiation

Microwave radiations can shorten the reaction times and increase the yield and purity of the products compared with the conventional heating methods. When solvents of high vapor pressures are employed in microwave (MW) irradiation treatment, there is a risk of explosions. Thus, the so-called dry media technique (i.e. in the absence of solvent) has been coupled to MW irradiation. The lack of use of organic solvents during the organic reactions gives rise intrinsically to cleaner, and more efficient and economic technology, increasing safety and decreasing the overall work needed for the process. In many cases, higher amounts of reactants can be treated and there is possibility of acting on the selectivity of the systems (Mingos and Baghurst, 1991).

1.4.1.3.7 Ionizing radiation

Ionizing radiation technology, an emerging AOP, characterized by the in situ

• − generation of highly oxidizing and reducing species (i.e., OH and eaq ) via equation (1.6)

(Spinks and Woods, 1990), has recently gained much attention because of the high efficiency, cost effectiveness and environmental compatibility (Bhatti et al., 2014).

Consequently, extensive research has been carried in the field of gamma irradiation for the destruction of toxic organic compounds including pesticides, herbicides, polychlorinated biphenyls (PCBs), methyl tert-butyl ether (MTBE), pharmaceuticals and dyes (Cooper et al., 2009, Basfar et al., 2007, Csay et al., 2012, Getoff, N., 1986, 1995,

Mincher et al., 1991, Rauf and Ashraf, 2009, Tahri et al., 2010). The end product of gamma radiation process is mostly carbon dioxide or other simple biodegradable and harmless end products.

• − • + H2O ---^-^--> OH(2.8) + eaq (2.7) + H(0.6) + H3O (3.2) + H2O2(0.72) +

− H2(0.45) + OHaq (0.5) (1.6) where the numbers in parenthesis denote the radiation yield or G-values at 10−7 s after irradiation. G-values are the number of species produced or destroyed per 100 eV of energy absorbed and are a means of expressing the efficiency of the radiolysis (Spinks and Woods, 1990).

− Hydrated electron (eaq ) is the most powerful reductant in aqueous solution (E˚ =

− 2.9 V) that reacts rapidly with many species having more reduction potentials (Buxton et al., 1988). An enhanced reactivity is observed when organic molecules contain electron withdrawing constituent such as halogen atoms, subsequently leading to the dechlorination of the organic compound (Buxton et al., 1988, Spinks and Woods, 1990).

Despite high affinity of the hydrated electron for reacting with chlorinated compounds, there is very limited study on the degradation of lindane, a typical chlorinated pesticide, by gamma radiation based AOPs (Mincher et al., 1991, Mohamed et al., 2009).

1.4.1.3.8 TiO2 photocatalysis

Titanium dioxide (TiO2) is one of the most promising photocatalysts because of its high efficiency, low cost, chemical inertness and photostability in nature (Hoffmann et al., 1995, Wilcoxon and Thurston, 1998). TiO2 photocatalysis is regarded as a promising technology for the treatment of water and wastewater organic pollutant (Lagunas-Allué et al., 2012, Malato et al., 2001, Ollis and Al-Ekabi, 1993). Around 40 years ago, Carey and co-workers reported the photocatalytic degradation of biphenyl and chlorobiphenyls in the presence of TiO2 (Carey et al., 1976). Since Carey’s invention, TiO2 photocatalysis was effectively applied for the destruction of large number of toxic compounds including

POPs (Hoffmann et al., 1995).

Several researchers have studied TiO2 photocatalytic degradation of lindane under

UV light (Dionysiou et al., 2000; Senthilnathan and Philip, 2010). Zaleska and co- workers reported 50% of lindane removal ([lindane]0 = 0.137 mM) in150 min, using TiO2

(anatase) supported on glass hollow microspheres (Zaleska et al., 2000). Dionysiou and co-workers reported 63% lindane removal in 170 min (([lindane]0 = 16.0 μM), using

TiO2 immobilized on a continuous flow rotating disc (Dionysiou et al., 2000).

Recent advances in the field of TiO2 photocatalysis introduce non-metal doping into

TiO2 that remarkably increases researchers’ interest because of a potential of the doped

TiO2 material to utilize low energy photon (visible light) for excitation of electron

(Hoffmann et al., 1995, Asahi et al., 2001). As a result, the restriction of UV light radiation for band gap excitation in TiO2 photocatalysis is removed, and consequently, this achievement has opened up various novel applications of TiO2, particularly in the sense of utilizing the most abundant solar light radiations (comprising ~ 45% visible

light) (Hoffmann et al., 1995, Chatterjee and Dasgupta, 2005). Developing new doped-

TiO2 photocatalyst is a rapidly growing research area and many studies have reported the synthesis, fabrication and characterization of the doped-TiO2 photocatalyst (Pelaez et al.,

2012). The potential application of visible light assisted doped-TiO2 photocatalysis for environmental remediation has been reported previously (Pelaez et al., 2012, Chatterjee and Dasgupta, 2005).

S-doping effectively achieves narrowing of the band gap of TiO2 or introduces localized mid-gap states in the band gap of TiO2 in a similar way as N-doping (Asahi et al., 2001). Sulfur appears to be one of promising non-metals for the synthesis of visible light-activated photocatalysts to decompose different contaminants (Umebayashi et al.,

2002). Sulfur was successfully inserted into TiO2 lattice via substitution of either oxygen

(O) or titanium (Ti) (Tachikawa et al., 2004). Umebayashi et al. reported that mixing of the S 3p states with the valence band (VB) caused an increase in the VB width, with a subsequent band gap narrowing in sulfur-doped TiO2 (S-TiO2), thereby shifting considerably the absorption edge to lower energy region (Umebayashi et al., 2002). Our research group has recently synthesized a nanostructured S-TiO2 photocatalyst through a sol–gel method, which is cost-effective and easy to operate (Han et al., 2011). Sol–gel method allows facile synthesis of photocatalysts with a high purity, stability and uniformity in their structural and physicochemical properties, at ambient conditions

(Mutuma et al., 2015).

The photocatalytic activity of TiO2-based photocatalysts for degrading organic pollutants can be increased in the presence of inorganic oxidants such as hydrogen

− 2− peroxide (H2O2), peroxymonosulfate (PMS, HSO5 ) or persulfate (PS, S2O8 ) (Andersen

et al., 2013, Chen et al., 2012, Malato et al., 1998). Compared to H2O2 or PS, PMS was found to be the most suitable oxidant for degradation of 2,4-dichlorophenol, Acid Orange

7 and other organic contaminants (Chen et al., 2012, Malato et al., 1998).

1.4.1.3.9 Sulfate radical based-AOPs

Compared to hydroxyl radical based AOPs (•OH-based AOPs) (represented by ozonation, UV/H2O2, photo-Fenton process and TiO2 photocatalysis), sulfate radical

•− based AOPs (SO4 -based AOPs) are relatively a new group of treatment technologies, which require more investigations on their potential applicability in treating the

•− contaminated water (Watts and Teel, 2006). SO4 can be generated by the activation of persulfate (PS) and peroxymonosulfate (PMS) with transition metals, elevated temperature or pH, and/or UV irradiation (Anipsitakis and Dionysiou, 2004a, 2004b).

•− • SO4 , with standard reduction potential, E˚= 2.5-3.1 V, is a stronger oxidant than OH

(i.e., E˚= 1.8-2.7 V), and is capable of rapid mineralization of recalcitrant pollutants at

•− neutral pH. The high oxidation efficiency of SO4 , high aqueous solubility, high stability with slow rate of the oxidants’ consumption, economically cheap and environmental friendly nature make this technology an attractive option in wastewater treatment and other environmental applications in recent years. Anipsitakis and Dionysiou (2003)

•− • reported that SO4 is more efficient than OH for the transformation of 2, 4- dichlorophenol, atrazine and naphthalene under certain conditions. Previous research

•− studies have demonstrated that SO4 -based AOPs employing PS as a precursor oxidant can effectively degrade lindane in water and soil (Cao et al., 2008, Usman et al., 2014).

However, no research studies have been found in the literature about the degradation of

•− lindane by SO4 -based AOPs employing peroxymonosulfate (PMS).

Peroxymonosulfate (available as a triple potassium salt with a commercial name of

® Oxone , 2KHSO5·KHSO4·K2SO4) is a highly versatile and an environmentally friendly oxidant (Kennedy and Stock, 1960). It has received great attention and application in water disinfection and decontamination (Anipsitakis et al., 2008). UV light and Fe2+ were

•− • chosen in this study to activate PMS, a process that can generate both SO4 and OH as shown below in reactions (1.7) and (1.8) (Anipsitakis and Dionysiou, 2004a, 2004b).

− •− • HSO5 + hν → SO4 + OH (Φ = 1.04) (1.7)

2+ − 3+ •− − 4 −1 −1 Fe + HSO5 → Fe + SO4 + OH (k = 3.0 × 10 M s ) (1.8)

1.5 Aims and objectives of the Present Work

 The main objective of this research work is to investigate the efficiency of several

emerging AOPs, i.e., gamma radiation, sulfate radical based-AOPs, and non-

metal doped TiO2 photocatalysis for the degradation of lindane in water.

2+ −  The efficiency of Fe /HSO5 based AOPs for the degradation of lindane will be

investigated under the influence of fluorescence light and UV light irradiation.

 The effects of water quality and process parameters, such as solution pH, initial

2+ − concentrations of lindane, Fe and HSO5 , presence of inorganic ions and natural

organic matters, will be studied.

−  The mineralization of lindane and decomposition of HSO5 under various AOPs

will be assessed.

• •− −  The second-order rate constant of lindane with OH, SO4 and eaq , generated in

the studied AOPs will be determined, using pulse radiolysis techniques or

competition kinetics model.

 Mechanism of lindane degradation in the selected AOPs will be investigated,

based on the reaction by-products, identified via GC/MS analysis.

 Dechlorination and mineralization efficiencies of various AOPs for degradation of

lindane will be studied, to assess the extent of water detoxification.

 A nano-structured non-metal doped S-TiO2 photocatalyst (e.g., S-TiO2) will be

synthesized by a sol-gel method using a self-assembly technique, and its

photocatalytic activity for the degradation of lindane under solar and visible light

irradiation will be investigated.

−  The effect of HSO5 on visible and solar light-assisted S-TiO2 photocatalysis of

lindane will be particularly studied.

 The degradation of trichlorobenzene (TCB), a typical reaction by-product of

lindane, will also be investigated.

 Finally, the assessment of lindane residues in the surface water samples, in

different regions of district Swabi, Khyber Pakhtoonkhwa (Pakistan) will be

carried out.

2. EXPERIMENTAL

Different types of materials, chemicals and instruments used during the present study along with procedure adopted and calibration of instruments are explained in this chapter.

2.1 Materials

Lindane (1,2,3,4,5,6-Hexachlorocyclohexane, γ-isomer; 97% purity) was obtained from Sigma Aldrich (UK). Hydrogen peroxide (50%, v/v) and ferrous sulfate (98%) used for photo-Fenton reagents, were purchased from Fisher Scientific (USA). Potassium peroxymonosulfate (PMS), available with commercial name oxone®

(2KHSO5.KHSO4.K2SO4, 95%) and sodium persulfate (Na2S2O8, 98%) were purchased from Sigma Aldrich. Inorganic salts e.g., sodium nitrate (NaNO3), sodium nitrite

(NaNO2), sodium sulfate (Na2SO4), sodium bicarbonate (NaHCO3), potassium chloride

(KCl) and potassium carbonate (K2CO3) were of analytical reagent grade quality and obtained from Fisher Scientific. Organic solvents and radical scavengers, such as ethanol, methanol, acetone, acetonitrile, tert-butanol (t-BuOH) and iso-propanol (i-PrOH) were all

HPLC grade and purchased from Merck (Germany). Humic acid (HA) standard was obtained from the International Humic Substances Society (IHSS, University of

Minnesota, St. Paul, MN, USA) and used as representative of natural organic matters

(NOMs). Sodium hydroxide and perchloric acid, used for pH adjustments, were purchased from Riedel-de Haën (Germany). Phosphate buffer solutions (pH 4.0 and pH

7.0) were used for pH meter calibration and were purchased from Scharlau (Spain).

Polyoxyethylene (80) sorbitan monooleate (Tween 80), isopropyl alcohol (i-PrOH,

99.7%), titanium (IV) isopropoxide (TTIP, 97%) and sulfuric acid (H2SO4, 95–98%) were used as precursors in the synthesis of TiO2 and S-TiO2 films and obtained from

Sigma–Aldrich (USA). Oxalic, tartaric, acetic and formic acids were obtained from Fluka and these acids were used as standard compounds for identification of reaction intermediates and products. Ultra pure nitrogen (carrier gas for GC, purity 99.999%) and ultra pure Helium (make up gas for GC, purity 99.999%) were obtained from Wright

Brothers Inc, PA (USA). Nitrous oxide (Medical grade) and oxygen gas (Commercial grade) used as dissolved gases were purchased locally in Peshawar. Ultra pure water

(resistivity of 18.2 MΩ.cm) was obtained from Milli-Q coupled with MILLIPORE Elix®

5 UV water purification system.

All the reagents were used as received without further purification.

2.2 Sample Collection

Surface water samples were collected from 18 different places in district Swabi. The sampling site is shown in Figure 2.1. The Swabi district is located on the bank of the

Indus River near Islamabad. Swabi district comprises a vast cultivation region covering a total area of about 1,550 square kilometers and having a population of around 1,027,000.

The major sources of irrigation water in Swabi district are the two main canals, i.e. the upper Swat canal and the Stefa canal. For drinking purposes and other house hold applications, people usually take water from the wells and tube wells, in district Swabi.

The samples were collected from the agriculture fields as well as from the canals. The samples SS1, SS2 (Shewa), SS4, SS5 (Kalu Khan) and SS8 and SS10 (Firdous abad) were taken from canal water whereas the remaining samples were collected from agriculture fields. The samples were stored in clean-washed plastic bottles and were analyzed in the Radiation and Environmental Chemistry laboratories (RECLs) of the

National Centre of Excellence in Physical Chemistry, University of Peshawar. Prior to

SPME technique, the samples were filtered through 0.45 μm filter paper in order to remove any type of particulate matter, if present.

Shewa Swabi SS1 SS2

SS6 SS7 Kalu Khan SS3 Adina Ismaila SS5

SS4 SS9 SS8 Terwatu Firdous SS11 Abad Yar Hussain SS12 SS10 Dobian Dagi

SS13 Yaqubi Sard Cheena SS15 SS14 SS16

SS17 Sudher

Ayuab Khan Kalay Sabdar Abad SS18

Figure 2.1. Sampling cite of district Swabi, KPK, Pakistan.

2.3 Preparation of Solutions

Stock solution of lindane (17.15 μM) was prepared by stirring the required weight of solid lindane in water for 24 hours at room temperature (T = 25 ± 1 ˚C). Standard solutions of lindane were prepared by dilution of the stock solution to the desired level.

° PMS (10 mM) and FeSO4.7H2O (5mM) stock solutions were prepared and stored at 4 C.

The pH of FeSO4.7H2O solution was set at 3.0 using 0.01M H2SO4. All the solutions were prepared in ultra pure Milli-Q grade water.

pH of the solutions was measured by a HANNA HI 9124 (USA) pH meter using glass electrode calibrated with standard buffers at pH 4.0 and 7.0. The calibrated pH meter was used for adjusting the pH of the test solutions.

2.4 Extraction Technique

Solid Phase Micro Extraction (SPME) technique was employed for extraction of lindane from aqueous solution. The SPME fibre needle (DVB/CAR/PMDS type;

Supelco, USA) consisted of a 10 mm long and 100 μm thick fused silica fibers, coated with 95 μm thick polydimethylsiloxane (PMDS). SPME technique is based on the equilibration of the analyte between an aqueous phase and an organic polymer coated onto a fused silica fiber. On injection in GC column, the analyte is thermally desorbed into the inlet and sent by mobile phase into the detector device. The sample injection was performed using a CTC Analytics CombiPAL autosampler connected with the GC. In some cases, SPME manual syringe (Supelco, Bellefonte, USA) was used for the injection purpose. The time of adsorption for lindane onto the SPME fibre was 2 min, while desorption time inside the inlet was 1 min.

2.5 Reactor Design

2.5.1 Gamma rays reactor

Ionizing radiation experiments were performed using a Cobalt-60 gamma-rays source (model Issledovatel (USSR)) available in the Food and Soil Division (FSD) of the

Nuclear Institute for Food and Agriculture (NIFA), Peshawar (Pakistan). The source consists of twelve cylindrical Co-60 rods, which are arranged in a circle (internal diameter 0.2 m). The height of each rod is 2.1 x 10–2 m and diameter is 1.1 x 10–2 m. The gamma radiation source is protected externally by using a thick lead shield. The picture of the source is shown in Figure 2.2.

Figure 2.2. Cobalt-60 gamma rays source at NIFA.

2.5.2 UV reactor

UV degradation experiments were performed in a bench scale photoreactor consisting of a Pyrex glass Petri dish (100 mm diameter x 15 mm height) with a quartz cover. A 10 mL of lindane solution (C0 = 3.43 μM) was taken in the reactor vessel and it was irradiated from the top with two 15 W low pressure mercury lamps (UV-C lamp from Cole-Parmer, USA) that emitted light primarily at a wavelength of 253.7 nm. The reactor vessel was held on the top of a magnetic stir plate. A Teflon coated magnetic stir bar was used to mix the solution during irradiation. The reactor vessel was sealed with

Parafilm and cooled with a fan to prevent evaporation and maintain a constant temperature (T = 25 ± 1 °C). The mercury lamps were turned on 30 minutes prior to experiments for uniformity of the UV fluence. At selected times after initiating photolysis; a 150 μL sample was removed from the reaction vessel and transferred to a

200 μL glass insert held inside the HPLC vial. Any radicals within the sample were immediately quenched by adding 50 μL methanol and samples were analyzed on Agilent

6890 GC interfaced to an Agilent 5975 mass selective detector. Control experiments were performed on the solutions that were (i) not photolysed and (ii) without oxidant.

2.5.3 UV-Vis reactor

A borosilicate glass dish (dia. 100 mm) was used as a photoreactor for photocatalytic experiments. The sulfur doped TiO2 (S-TiO2) films were placed in the photoreactor containing lindane solution at desired concentrations. The films were washed with MilliQ water and then dried under an infrared lamp before the photocatalytic experiments. Two 15 W fluorescent lamps (Cole-Parmer) were used as a visible light source. For visible light irradiation, a UV block filter (UV420, Opticology)

was mounted under the light source and the light intensity determined by a broadband radiant power meter (Newport Corporation) was found to be 9.05 × 10−5W cm−2. The output spectrum after the filter has been reported by Han and coworkers (Han et al.,

2011). A UV-Vis light irradiation assembly is shown in Figure 2.3. A 0.2 mL sample was withdrawn from the reaction mixture at selected times (i.e. 0, 1, 2, 4, and 6 h). The concentration of lindane in the samples was quantified using Gas chromatography-mass spectrometry (Agilent Series 6890) with a HP-5MS capillary column (30 m × 0.25 mm ×

0.25 μm). The photocatalytic experiments were performed in an Advanced

SterilChemgardIII ClassII (Baker) biological cabinet for health and safety reasons.

Figure 2.3. Batch reactor used for UV-Vis radiation treatment.

2.5.4 Solar reactor

Simulated solar light experiments were performed in a bench scale photoreactor consisting of a Pyrex glass Petri dish (100 mm diameter x 15 mm height) with a quartz cover. The reactor vessel was sealed with Parafilm and cooled with a fan to prevent evaporation and maintain a constant temperature (T = 27 ± 1 ˚C). A 10 mL of lindane solution was taken in the reactor vessel and it was irradiated from the top with 300 W

Xenon lamp (67005, Newport, Oriel Instruments) which provided a spatially averaged light intensity of 47.1 mWcm−2 obtained from a broadband radiometer (Newport

Corporation). The broad spectrum of the lamp light was transformed to mimicking solar light radiation by introducing appropriate filters. The first filter, an Air Mass 1.5 Global

Filter (Newport Corporation), attenuated the irradiation to simulate sunlight corresponding to a 48.2º angle of incidence. The second filter was an FSQ-KG5 filter

(Newport Corporation), which is a heat absorbing filter. The output spectrum of the filter has been reported by J. Andersen and coworkers, demonstrating adequate simulation of solar radiation (Andersen et al., 2013).

2.5.5 Tube-light reactor

Tube-light radiation experiments were conducted under the typical room light conditions that existed in the laboratory having dimensions: width × length × height = 15

× 20 × 15 = 4500 feet3 and equipped with sixteen ordinary Philips lightening tubes. A typical reactor for the tube-light experiments consisted of a 40 mL clear amber glass with

PTFE/Silicone screw caps. A 10 mL of lindane solution was transferred to the reactor vessel after addition of the desired amount of oxidant (PMS) and catalyst (Fe2+). The reaction mixture inside the reactor vessel was kept on constant mixing in a rotator (at 50

cycle/min) to ensure the homogeneity of the solution. The experiments were performed at room temperature (T = 23±2 °C). At selected time intervals, 150 μL samples were removed and transferred to a 200 μL glass insert kept inside the HPLC vials and quenched with 50 μL of methanol.

2.5.6 Dark reactor

The same experimental conditions as mentioned for tube-light experiments

(section 2.4.5) also existed here, except that the tube-lights were turned off during the reaction period and the clear amber vials were replaced with dark glass vials in case of dark experiments. After removal from the reaction vessel, the samples inside the HPLC vials were covered by aluminium foil to protect from the effect of light before the analysis is done.

2.6 Calibration of Radiation Sources

2.6.1 Calibration of irradiation source

A Fricke dosimetry solution was used for calibration of the irradiation source

(Sehested, 1970). Fricke dosimetry solution was prepared by dissolving 22.2 mL of 98%

H2SO4, 0.278 gm of FeSO4.7H2O and 0.058 gm of NaCl in one litre of milli-Q water.

NaCl was added to remove the effects of dissolved organic impurities, if present in the solution (Spinks and Woods, 1990).

The solution was purged with oxygen gas (O2) for 20 min to make it O2-saturated prior to irradiation. The main reaction of Fricke dosimeter is oxidation of ferrous ion

(Fe2+) to ferric ion (Fe3+) under gamma irradiation in the presence of oxygen ((Spinks and

Woods, 1990). A 15 mL Fricke dosimetry solution in 20 mL Pyrex glass tube was kept inside the gamma radiation field at ambient temperature (25±2 °C). After selected time

intervals, the samples were removed and change in absorbance of the solution was measured using UV / Vis spectrophotometer at wavelength of 304 nm using unirradiated solution in the reference beam. All the experiments were carried out in triplicate and a mean value for the change in absorbance (∆OD) was obtained. The change in absorbance

(∆OD) was plotted versus irradiation time (see typical calibration plot in Figure 2.4) and slopes of the plots ((∆OD/min) were calculated. The dose rate (D•) was determined by using equation 2.1 (Sehested, 1970);

N × ΔOD/min × 100 D =A eV/g (2.1) ε×ρ×1000×G(Fe3+ )

D• = Absorbed dose rate

23 NA = Avogadro number (6.022 × 10 )

(∆OD/min) = Difference in absorbance between the irradiated and non-irradiated samples per min (slope of the calibration plots).

−1 −1  x = molar extinction co-efficient of ferric ion at 304 nm 2205 M cm .

3  = Density of the Fricke dosimeteric solution (1.024 g/cm )

Radiation yield ( G -value) of Fe3+ for gamma radiolysis of Fricke solution = 15.6

3+ By putting the values of NA, (∆OD/min), εx , ρ , and G(Fe ) in equation (2.1), we get

6.0223x 1023  100 ΔOD/min D  eV/g (2.2) 2205 1.024  103  15.6

D• = 1.7099 x 1018 × (∆OD/min) eV/g (2.3)

Since 1 eV/g = 1.6022 x 10−14 rad

So, D• = 1.7099 × 1018 × 1.6022 × 10−14 × ∆OD rad.

D• = 2.7396 × 104 × ∆OD rad min−1 (2.4)

Since 1 rad. = 10−2 Gy

D• = 2.7396 × 102 × ∆OD Gy min−1 (2.5)

∆OD = slope (min−1)

∆OD/min = slope

Putting the value of ∆OD in equation (2.5), the values of D• under different experimental conditions were determined. The dose rate inside the cylindrical sample compartment of gamma ray source can be reduced by using different thickness metallic containers as given below.

a. Open container

The highest dose rate can be obtained in open container (using no metallic shielding inside sample compartment). An aqueous acidic solution of ferrous sulfate was irradiated in an open container of the Co-60 source and the slope obtained was determined, as shown in Figure 2.4. The value of slope (0.0168) was put into equation

(2.5) and dose rate was determined as shown below

D• = 2.7396 × 102 × 0.0168 Gy/min = 4.603 Gy/min

D• = 4.603 × 60 Gy/hr = 276 Gy/hr (2.6)

b. Brass container

In this case, Fricke dosimetry solution was irradiated in a brass container instead of the open container. Figure 2.4 shows the plot and value of the slope (0.0121) was used in equation (2.5) to determine the dose rate.

D• = 2.7396 × 102 × 0.0121 Gy/min = 3.315 Gy/min

D• = 3.315 × 60 Gy/hr = 199 Gy/hr (2.7)

c. Brass and iron container

By irradiating solution of ferrous sulfate in an iron container adjusted inside a brass container, the slope obtained is shown in Figure 2.4. The value of the slope

(0.0055) was inserted in equation (2.5) and the dose rate was determined as below.

D• = 2.7396 × 102 × 0.0055 Gy/min = 1.507 Gy/min

D• = 1.507 × 60 Gy/hr = 90 Gy/hr (2.8)

2.6.2 Calibration of UV radiation source

Three different calibration methods; iodide/iodate actinometry (Rahn, 1997), ferrioxalate actinometry (Murov et al., 1993; Goldstein and Rabani, 2008), and a calibrated digital radiometer (Model IL 1700, XRD (XRL) 140T254 low profile germicidal probe, International Light, Co., Newburyport, MA) were employed to determine the average UV fluence rate (mW cm−2) of the UV radiation source.

The aqueous solution used for iodide/iodate actinometery consisted of 0.6 M KI and 0.1 M KIO3 in a 0.01 M sodium tetraborate decahydrate (Na2B4O7.10H2O) buffer solution. The iodide/iodate actinometric experiments were performed within four hours after the solutions were prepared. In case of iodide/iodate actinometery, the photoproduct

 is tri iodide ion (I3 ) (see reaction 2.9) that is highly photosensitive in the UV range and can be accurately quantified at  = 352 nm (with molar extinction coefficient of 352 =

27,600 M−1 cm−1). The quantum yield of this actinometer is  = 0.73 at 254 nm (Rahn,

1997).

The overall photochemical reaction is:

    8 I + IO3 + 3 H2O + h  3 I3 + 6 OH (2.9)

The absorbance of the actinometric solution was measured with spectrophotometer soon after preparation at wavelengths of 300 nm and 352 nm. The absorbance of the actinometric solution measured at 352 nm before irradiation was used as a blank

(A352(blank)). The solution was then irradiated in a 10 mL Petri dish with UV-C light source for time intervals of 5, 10, 15, 20, and 25 min in triplicate and absorbance was measured after each reading (A352(sample)).

The following formula was used for calculation of UV fluence rate.

[A (sample )  A (blank )] E  23.969 352 352  V (mL) (mW cm2) (2.10) Area(cm 2 ) Exposure time(s)

Where E: light intensity,

Area: area of the Petri dish in cm2,

Exposure time: the time of sample illumination in s, and

V: the volume of sample solution in mL.

The above equation can be written into the following form as:

2 A352(sample) − A352(blank) = [(Area(cm )/(23.969  V (mL)) ) E]  Exposure time(s) (2.11)

In the present study the volume was 10 mL and area of the Petri dish was 19.64 cm2, thus:

A352(sample) − A352(blank) = (0.0819 E)  Exposure time(s) (2.12)

A plot of [A352(sample) − A352(blank)] versus Exposure time (s) gave a straight line with slope equal to 0.0084, as shown in Figure 2.5. Dividing the slope by 0.0819, we can get the light intensity in mW cm–2, which was 0.10 mW cm–2..

E = 0.0084/0.0819 = 0.10 mW cm–2.

The source was also calibrated with ferrioxalate actinometry and calibrated digital radiometer, whose details has been given elsewhere in literature (Murov et al., 1993;

Goldstein and Rabani, 2008). The results of the three measurements were consistent with each other. The calibrated radiometer was used to check the fluence rate of the UV source every day before use.

2.7 Qualitative/Quantitative Analysis of Lindane and other Products

2.7.1 Gas chromatography with electron capture detector (GC/ECD)

The degradation of lindane was monitored by gas chromatography (GC, Agilent

6890N) equipped with HP-5 capillary column (30 m × 0.32 mm × 0.25 μm, J&W

Scientific) using a Ni63 electron captured detector (ECD). The conditions used for the analysis of lindane by the GC/ECD are presented in Table 2.1. Lindane gave a sharp peak at a retention time of 9.033 min, under the given experimental conditions. A typical chromatogram of lindane on GC/ECD system is given in Figure 2.6.

Quantification of lindane was carried out by peak area measurements based on external standards calibration plot. A typical calibration plot for lindane on GC/ECD is given in Figure 2.7, with linear correlation coefficient (R2) of 0.9995. The limit of detection (LOD) at signal-to-noise ratio of 3 (S/N = 3) was 0.008 μM and the limit of quantification (LOQ) at signal-to-noise ratio of 10 (S/N = 10) was 0.037 μM. All the analyses were performed in triplicate and the mean value was used for results.

2.7.2 Gas chromatography-mass spectrometry (GC/MS)

In order to identify the by-products generated during the degradation of lindane,

GC/MS analysis was performed employing a HP 6890 series GC equipped with a HP

5973 mass spectrometry. Separation of the sample components was achieved using HP-

5MS capillary column (30 m × 0.25 mm × 0.25 μm). The conditions used for the analysis of lindane on the GC/MS system are presented in Table 2.2. Mass spectra were obtained by the electron-impact mode (EI) at 70 eV, using scan mode (50-800 m/z) under the following conditions: pressure: 7.63 psi, purge flow: 26.5 mL.min-1, purge time: 1 min.

GC/MS information was matched with NIST mass spectra library for identification of unknown compounds.

2.7.3 Ion chromatography (IC)

The chloride ion (Cl-) produced during degradation of lindane was analyzed with

Metrohm, 800 series Ion Chromatograph (IC) equipped with anion self-regenerating suppressor, a dual-piston pump, a DS6 conductivity detector and an IonPac A Supp 4 separating column (250 mm × 4 mm). An IonPac AG19 guard column (250 mm × 4 mm) was inserted before separating column for entrapping the impurities. A mixture of 1.8 mM sodium carbonate and 1.8 mM sodium bicarbonate solution (50/50 by volume) was used as eluent, while 50 mM sulfuric acid was added as a regenerating agent. The flow rate was 1 mL min-1 and the injection volume was 500 μL. The eluents were sonicated for

15 min and filtered through 0.45 μm cellulose acetate filters (Sartorius, Ministar) before use, to prevent the effects of air bubbles and other impurities.

The same IC system was used for determination of aliphatic acids, as degradation by-products such as oxalic, acetic and formic acids, but the gradient was changed (5 mM from 0 to 10 min and then increasing to 10 mM from 10 to 40 min). Before injection in

IC, the sample was filtered through 0.45 μm cellulose acetate filters (Sartorius, Ministar) for removal of impurities.

2.7.4 High performance liquid chromatography (HPLC)

The concentration of 2-chlorophenol (used as reference compound in the competition kinetics study) was monitored by high performance liquid chromatography

(HPLC) using an Agilent 1200 series equipped with a UV detector. The separation was achieved on reversed phase Eclipse XDB-C18 column (4.6 mm × 150 mm, 5 μm). The column was thermostated at 30 ºC. The injection volume was 20 μL. The column was eluted with a mixture of water: acetonitrile at 50:50 (v/v) with a flow rate of 1.0 mL/min.

2.7.5 Total organic carbon (TOC) analyzer

Total organic carbon (TOC) of the treated and untreated samples was monitored using a Shimadzu VCSH-TOC analyzer equipped with an ASI-V autosampler.

Calibration of the TOC analyzer was performed with potassium hydrogen phthalate and sodium hydrogen carbonate, used as standards for measuring total carbon and inorganic carbon, respectively. The difference between total carbon and inorganic carbon gives

TOC of the sample. When the lindane degradation experiments were conducted at concentration of 3.43 μM, the TOC experiments were performed at concentration of 17.5

μM, as the lower concentration of lindane cannot work with TOC.

2.7.6 UV spectrophotometer

The concentration of PMS residue was determined using HP 8452A UV–vis spectrophotometer, according to a reported method (Liang et al., 2008). The spectrophotometer covered a wavelength range from 190 to 820 nm and was equipped with a diode array detector and a deuterium lamp. PMS concentrations were determined by comparing absorbances to a standard calibration plot, constructed from data obtained with known concentration of PMS processed under the same experimental conditions. In

the PMS concentrations range of 0 - 2.0 mM, a plot of absorbance Vs. PMS concentration generated a straight line with R2 = 0.990. A typical calibration plot is shown in Figure

2.8.

2.8 Synthesis of Sulfur Doped TiO2 Photocatalyst

Sulfur doped TiO2 photocatalyst was synthesized via a modified sol-gel method explained elsewhere in literature and briefly described here (Han et al., 2011). A nonionic surfactant polyoxyethylene (80) sorbitan monooleate (Tween 80, Sigma–Aldrich) was used as a pore directing agent. The surfactant (Tween 80) was dissolved in iso-propyl alcohol (i-PrOH, 99.8%, Pharmco) and titanium (IV) isopropoxide (TTIP, 97%, Sigma–

Aldrich) was added as an alkoxide precursor. Finally, sulfuric acid (H2SO4, 95–98%,

Pharmco) was added as a sulfur precursor and reagent for in situ formation of water. The final solution obtained was somewhat yellow like, transparent, homogeneous and stable after stirring for 24 h at room temperature. The Tween 80:i-PrOH:TTIP:H2SO4 molar ratio employed for the preparation of the sol was 1:45:1:1. As reference, TiO2 films were synthesized following the same procedure but the surfactant was excluded and sulfuric acid was replaced with acetic acid at the same molar ratio.

The resulting sol–gel was immobilized onto plain borosilicate glass micro slides

(Gold Seal, 75 mm × 25 mm) by dip-coating method. Before dip coating, the slide surface was rinsed with Milli Q water followed by ethanol and dried under an infrared lamp. A self-made dip-coating apparatus with a speed controller device was used to dip in and pull out the substrate from the sol at a withdrawal rate of 12.5 ± 0.3 cm min−1 for a final effective surface area of 10 cm2. After coating, the film was dried with an infrared lamp for 20 min and calcined in a multi segment programmable furnace (Paragon Model

HT-22-D, Thermcraft Inc., Winston-Salem, NC) where the temperature was increased at a ramp rate of 900 ºC h−1 to 400 ˚C, maintained for 30 min to remove all organics and then allowed to cool down naturally. The dip coating and calcination process were repeated five times to obtain films with 5 layers (thickness: 1.02 ± 0.02 μm, total mass:

4.51 ± 0.18 mg). In addition, sulfur doped TiO2 particles from thick films were prepared to characterize the porosity and crystal structure, since it is very difficult to collect samples from the glass substrates due to extremely small amount of TiO2 in the thin films. For sulfur doped TiO2 particle characterization, the sol was dried on borosilicate glass dishes at 90 ºC for 6 h and then heat-treated at 400 ºC for 12 h using a multi- segment programmable high temperature furnace (Paragon Model HT-22-D, Thermcraft

Inc., Winston-Salem, NC) in order to remove all organics completely, resulting in the formation of thick films. The TiO2 particles were scratched up from thick films and grinded. Reference TiO2 films and particles were prepared following the same preparation and calcination processes.

0.40

Open container, R2 = 0.9965 0.35 Brass containor, R2 = 0.9968 Brass + Iron containor, R2 = 0.9963 0.30

0.25

0.20

0.15

0.10

0.05

Change in absorbance at 304 nm (a.u.) at nm in304 Change absorbance

0.00 2 4 6 8 10 12 14 16 18 20 22 Irradiation time (min)

Figure 2.4. Typical calibration plots of gamma irradiation source using Fricke dosimetry solution kept in different containers (a.u = arbitrary units).

1.8

1.6

1.4

1.2

1.0

0.8

0.6

Absorbance (a.u.) Absorbance

0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of UV photolysis (s)

Figure 2.5. A typical calibration plot for UV radiation source at λmax = 352 nm, by iodide/iodate actinometry.

−1 Figure 2.6. A typical chromatogram of lindane ([C]0 = 10 µgL ) using GC/ECD system (The GC/ECD conditions are given in Table 2.1).

1800

1600

2 1400 R = 0.9995

1200

1000

800

Peak area (a.u.) Peak 600

400

200

0 0 5 10 15 20 25 Concentration of lindane (M)

Figure 2.7. A typical calibration plot for lindane measurement on GC/ECD system.

3.5

3.0

2.5

2.0

1.5

Absorbance (a.u.) Absorbance 1.0

0.5

0.0 0.0 0.5 1.0 1.5 2.0 2.5 Concentration of calibration solution (PMS, mM)

Figure 2.8. A typical calibration plot for PMS measurement on the UV-Vis

spectrophotometer at λmax = 352 nm.

Table 2.1 Operating conditions for GC/ECD analysis of lindane.

Items Conditions Injector split-state splitless Temperature (°C) 210 °C SPME extraction time (min) 2 SPME desorption time (min) 1 Oven Temperature program 50 °C (4 min) to 150 °C at 5 °C /min, 150 °C to 250 °C (5 min) at 8 °C /min Carrier gas Gas Nitrogen Flow rate (mL/min) 1.5 Detector µECD

Table 2.2 Operating conditions for GC/MS analysis of lindane.

Items Conditions

Injector split-state splitless Temperature (°C) 210 °C SPME extraction time (min) 2 SPME desorption time (min) 1 Oven Temperature program 50 °C (4 min) to 150 °C at 5 °C /min, 150 °C to 250 °C (5 min) at 8 °C /min Carrier gas Gas Helium Flow rate (mL/min) 1.5 Detector GC/MS Ion source temperature 230 °C Quadrapole temperatures 150 °C Electron energy 70 eV Scan range (m/z) 50-450

3. RESULTS AND DISCUSION

Lindane is highly toxic and persistence pesticide and incidence of lindane contaminations have been found in water and soil samples in different parts of the world.

Residue of lindane in the field water samples of district Swabi, which is agriculturally one of the most developed areas of the Khyber Pakhtoonkhwa (KPK), Pakistan, has been investigated.

3.1 Determination of Lindane in Field Water Samples of District Swabi, KP, Pakistan

3.1.1 Optimization of the GC/ECD method for lindane analysis

First of all, a specific method was developed for analysis of lindane in water on the GC/ECD system using SPME fibre. The optimized method parameters were given in

Section 2.3. The linearity of the detector response was studied in lindane concentration range of 0.02 to 20.00 µg/L. Within the given concentration range, the detector gave a linear response with R2 = 0.999. The calibration plot is shown in Figure 2.6 (chapter 2).

The limit of detection (LOD) determined at signal to noise ratio (S/N) of 3 was found to be 0.008 μg/L and the limit of quantification (LOQ) determined at signal to noise ratio of

10 was found to be 0.037 μg/L. The LOD, LOQ and R2 value of the GC/ECD are given in Table 3.1.

3.1.2 Precision, accuracy, reproducibility and relative recoveries of the method

The efficiency of the given method was checked by measuring the precision, accuracy and reproducibility of the system under the optimized conditions and the results are presented in Table 3.2. The accuracy of the optimized method was calculated by using equation (3.1) (Whitmire et al., 2010):

Average measured concentration Accuracy = × 100 (3.1) nominal concentration

The relative recoveries were also determined by using equation (3.1), but in this case field water was used as solvent for the preparation of lindane solution rather than

Milli-Q water, which was used in case of accuracy measurement. Precision (closeness of agreement between the replicate independent test results) was measured in terms of relative standard deviation (% RSD). It can be seen from Table 3.2 that very precise, highly accurate and reproducible results were obtained under both intra-day and inter-day conditions. The relative recoveries of the optimized method are shown in Table 3.3.

Table 3.1. Performance of SPME followed by GC/ECD for determination of lindane

Analyte Equation Correlation coefficient LOD (µg/L) LOQ (µg/L) (R2)

Lindane y = 67.02x + 2.75 0.9995 0.008 ± 0.0003 0.037 ± 0.002

Table. 3.2. Intra-day and inter-day precision, accuracy and reproducibility of the optimized method

Analyte Nominal Intra-day response Inter-day response

concentration measured Precision Accuracy measured Precision Accuracy

(μg/L) concentration (%RSD) (%) concentration (%) (%RSD)

(Mean ± SD) (Mean ± SD)

Lindane 2.5 2.64 ± 0.27 10.29 105 2.37 ± 0.28 11.91 94.78

in water 5.0 4.82 ± 0.25 5.15 96 5.03 ± 0.26 5.12 100.55

10.0 9.20 ± 0.71 7.68 92 9.27 ± 0.39 4.21 92.70

Table. 3.3. Comparison of relative recoveries (Accuracy) and relative standard deviation (Precision) of lindane by SPME-assisted

GC/µECD method in different field water samples.

Lindane

conc. SS1 SS3 SS6 SS11 SS14 SS15

(μg/L)

Relative Relative Relative Relative Relative Relative R.S.D. R.S.D. R.S.D. R.S.D. R.S.D. R.S.D. recovery recovery recovery recovery recovery recovery

(%) (%) (%) (%) (%) (%) (%) (%) (%) (%) (%) (%)

2.5 88.18 8.98 110.57 7.87 88.58 13.18 114 2.64 103.37 16.04 99 17.48

5.0 98.46 9.16 101.05 11.28 95.66 14.84 103 8.99 96.56 3.11 92 7.77

10.0 83.31 9.78 88.40 8.09 99.69 7.14 99 2.56 88.05 5.17 100 5.31

3.2 Application of the Optimized GC/ECD Method to Real Water Samples

The validated method was applied for analysis of 18 field water samples collected from different places of Swabi district. The results obtained are given in Table 3.4. It can be seen from Table 3.4 that out of eighteen (18) samples analyzed, thirteen (13) were found to be contaminated with lindane, while in remaining five (5) samples, no lindane was detected. Of the thirteen contaminated samples, seven samples were found to contain lindane concentration beyond the maximum acceptable level (MAL) for a single pesticide in surface water i.e., 1.0 µg/L (Lopez-Blanco et al., 2002), while the remaining six samples had lindane concentration below the MAL. It can also be seen from Table 3.4 that canal water (i.e. Samples SS1, SS2 (Shewa), SS4, SS5 (Kalu Khan) and SS8 and

SS10 (Firdous abad) contains more amount of lindane pesticide as compared to the field water. A comprehensive detail about the sample information is provided in section 2.2 and Figure 2.1 (chapter 2). The possible reason is that pesticide contaminated water of various agriculture fields is continuously added which can increase the concentration of lindane in the canal water. The presence of high concentrations of lindane in the water samples of district Swabi is mainly due to the use of these pesticides for various agriculture purposes in this area. The high concentration of lindane (above the maximum acceptable level) presents high risks for various pesticide related diseases in the inhabitants of this region.

After the first step of evaluating the residues of lindane in some of the field water samples of district Swabi, the next step of the current work was to develop a chemically advanced and environmental friendly technique for removal of lindane from aqueous solution. Gamma radiation-induced degradation of lindane was studied under various

experimental conditions and the effect of different parameters on the degradation kinetics was determined. To avoid the effect of the matrixes found in field water, the degradation experiments were carried out in ultra pure Milli-Q water. However, the effect of natural organic matter (NOM) and inorganic salts (present in field or surface water) on the removal efficiency was determined by mixing these substances into the Milli-Q water.

Finally, the absolute rate constant of hydroxyl radical with lindane was determined by using competition kinetics, while the absolute rate constant of hydrated electron with lindane was determined by pulse radiolysis technique. In addition to the gamma radiation technology, removal of lindane from water by several other innovative AOPs, such as

•− TiO2 photocatalysis, SO4 radical-based Fenton-like and photo-Fenton-like processes were also investigated.

Table 3.4. Distribution of lindane in the field water samples of district Swabi.

Sites Concentration (µg/L) Sites Concentration (µg/L)

SS1 3.31 SS10 4.03

SS2 3.77 SS11 ND

SS3 ND SS12 0.40

SS4 4.13 SS13 ND

SS5 3.54 SS14 0.06

SS6 ND SS15 ND

SS7 0.17 SS16 1.35

SS8 1.08 SS17 0.44

SS9 0.09 SS18 0.11

ND: Not detected

3.3 Gamma Radiation-induced Degradation of Lindane in Water

An aqueous lindane solution (C0 = 3.43 µM) was irradiated by gamma rays and the concentration ratio (concentration at time ‘t’/ initial concentration = C/C0) was plotted against absorbed dose. The results of the gamma radiolytic decay of lindane are shown in

Figure 3.1. As can be seen, the concentration of lindane decreases with increase of gamma radiation doses which ultimately led to 97% lindane removal at an absorbed dose of 2000 Gy.

Radiolysis of water results in the generation of highly reactive radical species,

− • • such as e aq, OH and H (equation (1.6) that attack and destroy the organic pollutants

(Spinks and Woods, 1990). Figure 3.1 shows that the lindane degradation rate is fast in the beginning, while it slows down with increasing accumulated radiation dose. This behavior can be explained on the basis of the competition for the pollutant (lindane) between the reactive radicals, which increases with increasing radiation dose, thus resulting in decreased removal efficiency (Yu et al., 2008). Another possible explanation is the competition of the intermediate by-products with the lindane molecules for the reactive radicals which increase with the increasing accumulated radiation dose. Such a decrease in the degradation rate with the increasing radiation dose was observed in the radiolytic degradation of several other pollutants as well (Lin et al., 1995; Basfar et al.,

− • 2005a; Yu et al., 2008). The radical–radical recombination reactions, including e aq, OH and H• that increase with the increasing radiation dose rate (Lin et al., 1995) can be another possible reason for the decreasing removal rate at increasing accumulated radiation dose. Some of the major reactions along with bimolecular reaction rate

constants (mol−1 s−1) that can take place in the gamma irradiated aqueous solution are given below (Buxton et al., 1988, Spinks and Woods, 1990):

• − − 10 −1 −1 OH + eaq → HO (k = 3.0 × 10 M s ) (3.2)

• • 9 −1 −1 OH + H → H2O (k = 7.0 × 10 M s ) (3.3)

• • 10 −1 −1 OH + OH → H2O2 (k = 5.5 × 10 M s ) (3.4)

− − − 10 −1 −1 eaq + eaq → H2 + 2HO (k = 1.1 × 10 M s ) (3.5)

• • 9 −1 −1 OH + H2O2 → HO2 + H2O, (k= 2.7 × 10 M s ) (3.6)

• • 10 −1 −1 OH + HO2 → H2O + O2, (k = 1 × 10 M s ) (3.7)

• − •− 9 −1 −1 OH + HO2 → H2O + O2 , (k= 7.5 × 10 M s ) (3.8)

3.3.1 Kinetic Studies of the Gamma Radiation-induced Degradation of Lindane

The exponential decrease in lindane concentration with the increasing radiation dose suggested pseudo-first order kinetics behavior for lindane decay, expressed by equation (3.9) (Spinks and Woods, 1990).

C = C0 exp (−kD) (3.9)

The modified version of equation (3.9) is given below:

−ln[C/C0] = k.D (3.10)

where C0 is the initial concentration of lindane; C is the residual concentration of lindane at any radiation dose; D is the absorbed dose; and k is the dose constant in

˗1 reciprocal dose units (Gy ). When −ln[C/C0] is plotted against the absorbed dose (D), a straight line is observed whose slope is equal to the dose constant, k. The value of dose constant, k can be affected by several experimental conditions, such as initial pollutant concentration, solution pH, the addition of radical scavengers and the molecular structure of the compound (Mincher et al., 2002; Lee and Jeong, 2009).

3.3.1.1 The effect of initial solute concentration

The effect of initial lindane concentration on the gamma radiation-induced degradation of lindane was studied in batch kinetic experiments. Initial lindane concentrations of 0.343, 0.686, 1.715 and 3.434 μM, were irradiated for absorbed doses up to 1000 Gy and the lindane decomposition results are shown in Figure 3.2a.

The normalized lindane concentration (C/C0) decreased with increasing absorbed dose. For absorbed dose of 1000 Gy, the removal efficiencies for initial aqueous lindane concentrations of 0.343, 0.686, 1.715 and 3.434 μM were 95, 94, 91 and 83%, respectively. The observed degradation dose constants of lindane at different initial concentrations are shown in Figure 3.2b and Table 3.5. The dose constant increased with decreasing initial concentrations of lindane. An increase in the degradation dose constant with decreasing initial concentrations of the pollutants was also observed during the gamma radiolysis of 2,4,6-trinitrotoluene (Lee and Lee, 2005), cefaclor (Yu et al., 2008) and alachlor (Choi et al., 2010).

1.0 4

3 2 0.8 R = 0.9994

)

0 2

0.6 -ln(C/C 1

(lindane)

0 0.4 0

C/C 0 500 1000 1500 2000 2500

Absoseorbed dose (Gy) 0.2

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

Figure 3.1. Radiation-induced degradation of lindane in N2-saturated aqueous solution.

Experimental conditions: [lindane]0 = 3.43 µM, pH = 6.8. In the insert is shown

the plots of −ln(C/C0) Vs absorbed dose for determining pseudo-first order kinetics.

Table 3.5. Change in dose constant (k) with increasing the initial concentration of lindane. −1 2 [Lindane]0 (µM) Dose constant, k (Gy ) R

0.343 0.0029 0.9958

0.686 0.0027 0.9954

1.715 0.0023 0.9988

3.434 0.0017 0.9895

(a)

1.0

C0 = 3.434 M

C0 = 1.715 M

0.8 C0 = 0.694 M

C0 = 0.343 M

0.6

(lindane)

0 0.4

C/C

0.2

0.0 0 200 400 600 800 1000 1200 Absrobed dose (Gy)

(b)

3.5 2 C0 = 0.343 M, R = 0.9958 2 3.0 C0 = 0.694 M, R = 0.9955 2 C0 = 1.715 M, R = 0.9988 C = 3.434 M, R2 = 0.9896 2.5 0

)

0 2.0

-ln(C/C 1.5

1.0

0.5

0.0 0 200 400 600 800 1000 1200 Absrobed dose (Gy)

Figure 3.2. (a) Removal of lindane at different initial concentration Vs irradiation doses and (b) dependency of the initial concentration on the dose constant. Experimental

conditions: pH = 6.8, N2-saturated.

3.3.1.2 Effect of pH

The efficiency of gamma irradiation process is greatly associated with the solution pH. The results of the gamma radiolytic degradation of lindane under three different pH conditions (4.0, 8.0 and 6.8) are shown in Figure 3.3a. The results show that the lindane removal efficiency is maximum at neutral pH, while it decreased under both acidic and basic conditions. At an absorbed dose of 2000 Gy, the degradation efficiency of lindane was 50, 97, and 80% when the solution pH was 4.0, 6.8 and 8.0, respectively.

The removal efficiencies of the radiation-induced degradation processes are dependent on the type and concentration of active species. The concentrations of the reactive species of water radiolysis vary with the solution pH (Spinks and Woods, 1990). In acidic medium,

− + the hydrated electron (eaq ) reacts with H (equation (3.11) (Zhang et al., 2007). In alkaline solution, •OH can react with −OH (equation (3.12)), thereby decreasing the concentration of •OH (Buxton et al., 1988, Guo et al., 2009). The equations (3.11) and

(3.12) are likely to lower the concentration of hydrated electron and •OH radicals resulting in reduced lindane removal efficiency under the acidic and basic conditions.

− + • 10 −1 −1 eaq + H → H k = 2.3 × 10 M s (3.11)

• − − 10 −1 −1 OH + HO → H2O + O k = 1.3 × 10 M s (3.12)

Lindane removal efficiency at different pH values fitted pseudo-first order kinetic model.

At solution pH of 4.0, 6.8 and 8.0, the observed degradation dose constants are 3.56 ×

10−4, 1.71 ×10−3 and 8.21 × 10−4 Gy−1, respectively. The observed degradation dose constants of lindane obtained from the plots of −ln(C/C0) vs. absorbed dose at different initial pH are shown in Figure 3.3b. It can be seen that the degradation dose constant decreased both under the acidic and basic pH conditions. The observed decrease in

− degradation efficiency at acidic pH is most likely due to the scavenging effects of eaq by the H+ ion under such conditions (Spinks and Woods, 1990, Huang et al., 2009).

Similarly, the standard reduction potential of •OH is significantly reduced at higher pH and this may be another factor leading to a decreased removal efficiency at higher pH.

(a)

1.0 pH = 4.0 pH = 8.0 pH = 6.8 0.8

0.6

(lindane)

0 0.4

C/C

0.2

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

(b)

pH = 6.8, R2 = 0.9853 pH = 8.0, R2 = 0.9755 pH = 4.0, R2 = 0.9555 3

)

0 2

-ln(C/C

500 1000 1500 2000 2500 Absorbed dose (Gy)

Figure 3.3. (a) Removal of lindane at different initial pH Vs irradiation doses, and (b) Dependency of the initial pH on the dose constant. Experimental conditions:

[lindane]0 = 3.43 µM, N2-saturated. 82

3.3.2 Scavenging effects on gamma radiation-induced degradation of aqueous lindane

Several kinds of reactions may be involved in the gamma radiolytic degradation of lindane in aqueous solution. Lindane may be degraded by reaction with reducing

− • species such as, aqueous electron (eaq ) and hydrogen radical (H ), or it may be oxidized by reaction with hydroxyl radicals (•OH). The radicals produced in the gamma radiolysis of water can be selectively scavenged by the addition of certain chemicals or gases that allow measurement of the role of a single radical with the substrate. In order to elucidate the role of a single reactive radical and to find out the overall mechanism of degradation, the effect of dissolved gases (such as nitrogen, air and nitrous oxide) on the lindane degradation was studied. The degradation studies were performed under a set of four different experimental conditions that included:

1. N2-saturated solution (control),

2. N2-saturated solution containing 60 mM i-PrOH (reductive conditions),

3. N2O-saturated solution (oxidative conditions) and

4. Air-saturated solution.

Figure 3.4a illustrates the ratio of lindane concentration (C/C0) as a function of irradiation dose, while Figure 3.4b shows the observed degradation dose constants under the above mentioned conditions. The results indicated that lindane degradation efficiency was maximum in the N2-saturated solution (control condition), while it was minimum in the N2O-saturated solution (oxidative conditions). The degradation efficiency had intermediate value in the aerated solution and N2-saturated solution containing i-PrOH.

3.3.2.1 Radiolysis of lindane in N2-saturated solution (control experiments):

In N2-saturated solutions, all of the originally produced reactive radicals i.e. hydrated electrons, hydroxyl radical and H-atoms (equation (1.6)) are present and able to react with lindane pesticide. It can be seen from Figure 3.4a and 3.4b that the highest degradation efficiency was observed under N2-saturated conditions, corresponding to

97% lindane removal at an absorbed dose of 2000 Gy. Under N2-saturated conditions, the stoichiometric ratio of the reactive radicals present in the aqueous solution is 46:44:10 for

• - • the OH, e aq and H radicals, respectively (Wasiewicz et al., 2006). The highest degradation efficiency under the N2-saturated conditions revealed that all the three major

− • • species of water radiolysis i.e. eaq , OH and H radicals can play significant role in the degradation of lindane.

3.3.2.2 Radiolysis of lindane in aerated solution (natural condition):

Under aerated conditions, the dissolved oxygen in water (ca. 2.5 × 10−4 M at 25

− ˚C) is able to scavenge a large fraction of the hydrated electrons (eaq ) and almost all of the hydrogen atoms (equations (3.13) and (3.14)) in the solution (Buxton et al., 1988). In

• •− aerated solutions, there are 46% OH and 54% O2 species reacting with substrate

(Wasiewicz et al., 2006).

− •− 10 −1 −1 eaq + O2 → O2 (k = 1.9 × 10 M s ) (3.13)

• + − 10 −1 −1 H + O2 → H + O2 (k = 1.2 × 10 M s ) (3.14)

3.3.2.3 Radiolysis of lindane in N2-saturated 60 M i-PrOH solution (reductive

conditions):

For selective monitoring of hydrated electrons, experiments were performed in a

N2-saturated 60 M i-PrOH solution. The function of the N2 pre-saturation is to remove

the dissolved oxygen (O2) from the solution. The dissolved oxygen acts as a strong

− • • scavenger of eaq , as shown in equation (3.14). i-PrOH scavenges OH and H to form the relatively inert radical as shown in equations (3.15) and (3.16) (Buxton et al., 1988).

Figure 3.4a shows that the presence of 60 mM i-PrOH had only a small negative effect on the degradation efficiency of lindane. These results revealed that •OH and H• radicals have minor contribution in the degradation of lindane.

• • 9 −1 −1 i-PrOH + OH → (CH3)2 COH + H2O (k = 1.9 × 10 M s ) (3.15)

• • 7 −1 −1 i-PrOH + H → (CH3)2 COH + H2 (k = 7.4 × 10 M s ) (3.16)

3.3.2.4 Radiolysis of lindane in N2O-saturated solution (oxidative conditions):

The role of hydroxyl radicals (•OH) was determined by pre-saturating the lindane solution with N2O gas that quantitatively converts the hydrated electrons to hydroxyl radical (equation (3.17)) (Spinks and Woods, 1990);

− − • 9 −1 −1 eaq + N2O + H2O → N2 + OH + OH (k = 9.0 × 10 M s ) (3.17)

• • Under the N2O-saturated conditions there is 90% OH and 10% H , while no hydrated electrons exist in the solution. Figure 3.4a shows that the lindane degradation efficiency was strongly reduced in the N2O-saturated solution. These results revealed that the hydroxyl radical, •OH participated only slightly in the degradation of lindane and

− major contribution was from reaction of eaq with lindane.

From the graph, it appears that for 2000 Gy of irradiation, about 20% degradation

• − • is from OH radicals, 60% from eaq and about 10% from H radicals.

(a)

1.0 N2O-satutrated Aerated

N2+ 60 mM i-PrOH 0.8 N2-sturated

0.6

(lindane)

0 0.4

C/C

0.2

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

(b)

2 N2-saturated , R = 0.9955 2 N2+ 60mM i-PrOH, R = 0.9958 2 Aerated R = 0.9984 R2 = 0.9893 3 N2O-saturated

)

-ln(C/C

0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

Fig. 3.4.(a) Radiation-induced degradation of aqueous lindane in (i) Aerated solution;

(ii) N2-saturated solution; (iii) N2O-saturated solution; (iv) N2-saturated solution containing 60 mM i-PrOH. (b) Dependency of the radical scavengers on the dose

constant. Experimental conditions: [lindane]0 = 3.43 µM, pH = 6.8. 86

3.3.3 Role of individual reactive species in lindane degradation

‒ • • In order to investigate the role of individual reactive species (i.e., eaq , OH and H radical) towards lindane degradation, the gradation experiments were performed under the following conditions:

1. N2-saturated solution with no scavenger (scavenger free),

2. N2-saturated solution containing 60 mM i-PrOH

2. N2-saturated solution containing 60 mM t-BuOH

‒ • • In the scavenger free solution, all the three species i.e., eaq , OH and H are operative.

In the i-PrOH containing solution, •OH and H• radicals are scavenged thus prevailing

− • condition for eaq only, while t-BuOH containing solution selectively removes OH

− • radicals, prevailing condition for eaq and H . Figure 3.5a represents the ratio of lindane concentration (C/C0) as a function of irradiation dose, while Figure 3.5b shows the observed degradation dose constants under the given experimental conditions.

The observed dose constant (k) in the given experimental conditions were calculated from the plots of – ln(C/C0) versus absorbed dose (Figure 3.5b) and the calculated values are given below (Table 3.6):

‒3 ‒1 kno scavenger = 1.76 × 10 Gy

‒3 ‒1 kt-BuOH = 1.47 × 10 Gy

‒3 ‒1 ki-PrOH = 1.18 × 10 Gy

‒ • From the above mentioned values of k, the observed dose constant ratio of eaq , OH and H• can be calculated as:

‒ ke aq : k•OH : kH• ki-PrOH : (kno scavenger ‒ kt-BuOH) : (kt-BuOH ‒ ki-PrOH)

1.18 × 10‒3 Gy‒1 : (1.76 ‒ 1.47) × 10‒3 Gy‒1 : (1.47 ‒ 1.18) × 10‒3

Gy‒1

1.18 × 10‒3 Gy‒1 : 0.29 × 10‒3 Gy‒1 : 0.29 × 10‒3 Gy‒1

4 : 1 : 1

‒ • • ‒ • • The ratio of k for eaq , OH and H showed that eaq play major role, while OH and H

play minor role in the gamma radiolytic lindane degradation. The increased reactivity of

‒ eaq in terms of lindane removal was also observed when the lindane degradation was

carried out in the presence of various radical scavengers.

The same experimental data could also be used to determine the quantum efficiency

(η) of each reactive radical with respect to degradation of lindane. By definition, the

quantum efficiency is “the number of molecules decomposed by a reactive species” (Liu

‒ et al., 2005). Thus, the quantum efficiency for hydrated electrons (ηeaq ) can be given as:

-  Total number of lindane molecules decomposed by eaq eaq - (3.18) total number of eaq produced

‒ ‒ The G-value of eaq (number of eaq produced per 100 eV of radiation energy

‒ absorbed) is used to calculate the total number of eaq produced during irradiation process

whereas the number of lindane molecules decomposed is calculated from the change in

‒ molar concentration multiplied by Avogadro’s number. Since the G-value for eaq is 2.7

‒ (Spinks and Woods, 1990), the total number of eaq produced during 1000 Gy (1 Gy =

6.24 × 1018 eV) irradiation will be:

6.24 1018 eV 2.7 species 1000 Gy    1.68  1020 Gy100 eV

Therefore;

106 6.02  10 23  ΔC i PrOH N A  eaq -  20  0.0101 (3.19) total number of eaq  10

where ΔCi-PrOH is the change in lindane concentration for 1000 Gy of radiation dose when

60 mM i-PrOH was used as radical scavenger.

Considering 2.8 and 0.6 as G-values for •OH and H• (Spinks and Woods, 1990), the total number of •OH and H• produced during 1000 Gy energy absorbed will be 1.75 ×

1020 and 3.74 × 1019, respectively. Therefore, the quantum efficiencies for •OH and H• can be calculated as:

6 23 ΔCNo scavenger ΔCt BuOH  N A 10  6.02  10     0.0092 (3.20)  total number of • OH 10 20

2.4 × 10-6 6.02 × 10 23 ΔCtiBuOH ΔC PrOH N A      0.0386 (3.21)  total number of H• 3.74 × 10 19

Where ΔCN0 scavenger and ΔCt-BuOH represent change in lindane concentration for 1000

Gy of radiation dose when 60 mM t-BuOH and no radical scavenger are used, respectively.

The results obtained revealed that the quantum efficiency of each reactive species is far less than unity. It means that only a small fraction of the reactive species of the water radiolysis is involved in lindane degradation, while most part of these reactive species is wasted. This can be explained on the basis of two plausible reasons. Firstly, the various reactive species of the water radiolysis react with each other, thus reducing the active concentrations of these species than the original concentrations (Lin et al., 1995). Species 89

with low G value will have low degree of self recombination reactions and the resulting high quantum efficiency as compared to the species with high G value. The self recombination reactions of these reactive species are shown in reactions (3.2) – (3.8).

Secondly, the various intermediate by-products of lindane degradation may compete with lindane for the reactive radicals which could possibly result in low quantum efficiency of these species for lindane degradation (Lin et al., 1995; Basfar et al., 2005b; Yu et al.,

‒ • • 2008). The quantum efficiency ratio for eaq , OH and H may be expressed as:

‒ ηeaq : η•OH : η•H = 0.0101 : 0.0092 : 0.0386 = 1 : 1 : 4

‒ Table 3.6. The dose constants of hydrated electron (ke aq), hydroxyl radical (k•OH) and the ‒ overall degradation dose constant (ke aq,•OH,H•), G-value (Species/100 eV), %

removal and D0.5 of lindane at 500 Gy. ______−1 Type of reaction G-value Removal (%) k (Gy ) D0.5 (Gy) D0.9 (Gy)

−• • −3 keaq OH,H 0.0431 65 1.77 x 10 393 1310

− ‒3 keaq 0.0298 45 1.14 × 10 608 2019

• ‒4 k OH 0.0073 11 2.24 × 10 3094 10279

(a)

i-PrOH 1.0 t-BuOH No scavenger

0.8

0.6

(lindane) 0 0.4

C/C

0.2

0.0

0 500 1000 1500 2000 Absorbed dose (Gy)

(b)

No scavenger 4 t-BuOH i-PrOH

)

0 2

-lin(C/C

0 500 1000 1500 2000 Absorbed dose (Gy)

Figure. 3.5. (a) Radiation-induced degradation of aqueous lindane in (i) N2-saturated

solution; (ii) N2-saturated + 60 mM i-PrOH, (iii) N2-saturated + 60 mM t-BuOH (b) dependency of 60 mM i-PrOH and t-BuOH on the dose constant.

Experimental conditions: [lindane]0 = 3.43 µM, pH = 6.8. 92

3.3.4 Dechlorination studies of lindane

In addition to the studies of lindane degradation by gamma irradiation, studies concerning its dechlorination were also conducted under the same experimental conditions. The dechlorination studies were based upon the concentration of chloride ion

(Cl−) released during the gamma irradiation of lindane solution. The results of the Cl− release with the increasing radiation dose are given in Figure 3.6. As can be seen from the

Figure 3.6, the concentration of Cl− increased with increasing radiation dose and reached

12.6 μM (61% of 3.43 × 6 = 20.58 µM Cl−) at an absorbed dose of 2000 Gy, corresponding to 97% lindane degradation. The results showed that the rate of dechlorination is high in the beginning but it slowed down with the passage of time. The enhanced dechlorination efficiency in the beginning is due to greater activity of the reactive species at start that decreased with time due to their reaction with the intermediates of lindane degradation. The slower dechlorination rate at later stage can also be attributed to the lower remaining concentration of lindane in solution. The lower dechlorination rate can also be attributed to the formation of compounds with low chlorine content that show slower dechlorination behavior compared to compounds with more chlorine content (Buxton et al., 1988). Similarly, stable intermediates can be formed which can resist dechlorination. Such a dechlorination trend is comparable to that observed in the radiolytic dechlorination of several other chlorinated organic compounds, such as 2-chlorophenol, chloroacetic acid and chloroform (Taghipour and Evans, 1997).

Further increasing the radiation dose up to 4000 Gy, the dechlorination efficiency reached 88%, while complete disappearance of lindane occurred at radiation dose of 2620

Gy (data not shown). It is also possible that some even more stable intermediates, like

chlorinated organic acids and aldehydes may be formed that can be degraded at still higher radiation doses. A slight difference between the theoretical and experimental values may be due to experimental and analytical errors.

Nevertheless, the results revealed that lindane readily undergoes dechlorination at mild gamma radiation doses. Thus, the dechlorination of lindane mainly occurs due to reaction of the hydrated electrons with lindane. The dechlorination from vicinal carbons as well as dehydrochlorination reaction introduces unsaturated bonds in the organic molecules. The resulting unsaturated molecule is generally more susceptible to oxidation via •OH attack.

From the results of the scavenging gases, it is clear that the lindane dechlorination followed similar path as observed for lindane degradation. Under the effect of various radical scavengers, the efficiency of dechlorination was decreased in the following order:

N2-saturated > Air-saturated > N2O-saturated solutions. Such a dechlorination trend can be explained on the basis of reactivity of reactive radicals. The hydrated electron is highly reactive in the N2-saturated solution, while it is totally absent from the N2O- saturated solution.

3.3.5 Effect of common inorganic salts on lindane degradation

Na2CO3, NaHCO3, NaNO3, NaCl and NaNO2 are among the inorganic anions that are commonly found in natural water resources and are likely to affect the efficiency of

−3 −1 the degradation process. Figure 3.7 shows the effect of 10 mol L Na2CO3, NaHCO3,

NaNO3, NaCl and NaNO2 on lindane degradation by gamma-rays irradiation. The results showed that the degradation efficiency is influenced by the addition of these ions. As shown in Figure 3.7, at a given dose, the reduction efficiency of lindane was somewhat 94

higher in the presence of sodium carbonate and sodium bicarbonate than that in the

2− + + − absence of these salts. Since CO3 could react with H3O , the inhibition of H3O on eaq

− was reduced and the concentration of eaq produced by gamma-ray irradiation increased

•− (Schmelling et al., 1998). The lindane can also be degraded by CO3 radicals, produced in equation 3.23. On the other hand, the addition of NaNO3 and NaNO2 resulted in

− reducing the removal efficiency since eaq was quickly scavenged by these ions (Singh

• and Kremers, 2002) and OH was scavenged by NaNO2 (Zhang et al., 2007). The

− • − − − scavenging of eaq and OH by NO3 , NO2 and Cl ions is shown in reactions 3.22-3.26

2− (Spinks and Woods, 1990, Buxton et al., 1988). A positive influence of CO3 -ions on the gamma radiolytic degradation of pollutants is consistent from other literature reports also

(Zhang et al., 2007). Similarly, a negative effect of NaNO3 and NaNO2 additives on the pollutants degradation has also been reported in literature (Mucka et al., 2003, Zhang et al., 2007)

• 2− •− − 8 −1 −1 OH + CO3 → CO3 + OH (k = 3.39 × 10 M s ) (3.22)

‒ ‒ 2‒ 9 −1 −1 NO2 + e aq → NO2 (k = 3.5 × 10 M s ) (3.23)

‒ • ‒ 10 −1 −1 NO2 + OH → NO2 + OH (k = 1.0 × 10 M s ) (3.24)

‒ ‒ 2− 9 −1 −1 NO3 + e aq → NO3 (k = 9.7 × 10 M s ) (3.25)

Cl‒ + •OH → ClOH•˗ (k = 4.3 × 109 M−1 s−1) (3.26)

4 14

 3 10



( 8 - [lindane] [Cl-] 2 6

2 Concentration of Cl

Concentration of lindane (

0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

Figure 3.6. Dechlorination and degradation studies of lindane versus radiation dose.

Experimental conditions: [lindane]0 = 3.43 µM, pH = 6.8, N2-saturated.

1.0 - NO2 NO3- - 0.8 Cl No additive (Control) - HCO3 CO 2- 0.6 3

(lindane)

0 0.4

C/C

0.2

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

2− − ˗ − − Fig. 3.7. Effect of CO3 , HCO3 , Cl , NO3 and NO2 ions on lindane degradation by 2− gamma-ray irradiation. Experimental conditions: [lindane]0 = 3.43 µM, [CO3 ]0

− ˗ − − = [HCO3 ]0 = [Cl ]0 = [NO3 ]0 = [NO2 ]0 = 1mM, N2-saturated. 96

3.3.6 Effect of H2O2 on lindane degradation

H2O2 is an inorganic oxide that is commonly used for advance oxidation treatment of water. It may directly react with pollutant but most often it generates highly reactive hydroxyl radicals on gamma radiolysis (reaction 3.27). However, it has high second-order rate constant for reaction with hydrated electron and hydrogen atom and may negatively affect those processes which are controlled by those species. The reaction of H2O2 with hydrated electron and hydrogen atom is shown by reactions (3.28) and (3.29),

• respectively (Buxton et al., 1988). When used in excess amount, H2O2 acts as an OH radical scavenger by reactions (3.30) (Spinks and Woods, 1990)

 rays  H22 O2 OH (3.27)

‒ • ‒ 10 −1 −1 H2O2 + eaq → OH + OH (k = 1.1 × 10 M s ) (3.28)

• • 7 −1 −1 H2O2 + H → OH + H2O (k = 9.00 × 10 M s ) (3.29)

• • 7 −1 −1 H2O2 + OH → H2O + HO2 (k = 2.7 × 10 M s ) (3.30)

Figure 3.8 shows the effect of three different concentration of H2O2 on the gamma-radiation induced degradation of lindane in water. As can be seen, the lindane removal efficiency was greatly reduced in the presence of H2O2. From the negative effect of H2O2 on the lindane removal efficiency, it is evident that hydrated electron is the major species involved in the gamma radiolytic lindane decay.

3.3.7 Effect of natural organic matter (NOM) and synthetic organic pollutants

Humic acids (HAs) are yellow- to black-colored macromolecular substances that constitute a considerable fraction of natural organic matter (NOM) in surface waters. The presence of NOMs generally has an effect on the rate of degradation of organic pollutants in water. Although irradiation of NOM produces some reactive species, its radical 97

scavenger effects are significant enough to potentially decrease the rate of pollutant degradation in practical applications (Nienow et al., 2008). To test this hypothesis, lindane degradation experiments were carried out in the presence of 1 mg/L humic acid.

Similarly, several kinds of synthetic organic pollutants can be found in the water sources in nature. Two common pollutants namely chloroform (CHCl3) and chlorophenol (2-CP) were selected as model synthetic pollutants and the effect of these pollutants on gamma radiation-induced degradation of lindane in aqueous solution was investigated. Figure 3.9 shows the effect of CHCl3, humic acid and 2-CP on gamma radiolytic lindane degradation in water. The results showed that all these organic substances had small negative effect on the degradation rate of lindane. The HA is a scavenger of •OH (Nienow et al., 2008), while its reaction with hydrated electron is not reported. CHCl3 and 2-CP

− • have high reaction rate constants with eaq and OH, however, results showed that lower concentrations of these pollutants do not significantly affect the lindane removal efficiency. From these results, it is obvious that the presence of humic acid, 2-CP and

CHCl3 in the low concentration ranges of real-world application do not affect lindane degradation significantly. The various reactions occurring in the gamma irradiated lindane solution containing 2-CP and CHCl3 can be the following (Haag and Yao, 1992,

Buxton et al., 1988):

‒ 10 −1 −1 CHCl3 + e aq → Products (k = 3.0 × 10 M s ) (3.31)

• • 7 −1 −1 CHCl3 + OH → CCl3 + H2O (k = 5.2 × 10 M s ) (3.32)

• 7 −1 −1 CHCl3 + H → Products (k = 1.2 × 10 M s ) (3.33)

• 10 −1 −1 2-ClC6H4OH + OH → Products (k = 1.2 × 10 M s ) (3.34)

‒ 9 −1 −1 2-ClC6H4OH + e aq → Products (k = 1.3 × 10 M s ) (3.35)

1.0

0.8

0.6

(lindane)

0 0.4

C/C

H2O2 = 0mM

H2O2 = 5 mM 0.2 H2O2 = 10mM

H2O2 = 20mM

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

Figure 3.8. Effect of concentration of H2O2 on lindane degradation by gamma-ray irradiation. Experimental conditions: [lindane]0 = 3.43 µM, pH = 6.8, N2- saturated.

1.0

CHCl3 Humic acid 0.8 2-CP Blank

0.6

(lindane)

0 0.4

C/C

0.2

0.0 0 500 1000 1500 2000 2500 Absorbed dose (Gy)

Figure 3.9. Effect of humic acid, chloroform and chlorophenol on lindane degradation

by gamma-ray irradiation. Experimental conditions: [lindane]0 = [2-CP] =

[CHCl3] = 3.43 µM, [humic acid] = 1 mg/L, N2-saturated. 99

− 3.3.8 Pulse radiolysis of lindane–Hydrated electron rate constants (eaq + lindane)

Pulse radiolysis can be used to determine fast reaction kinetics and spectra of short lived transients produced during radiolysis of aqueous solutions. The radicals produced in the electron pulse radiolysis of water can be selectively removed by the addition of certain scavengers, to allow measurement of the rate constant of a single radical with a substrate. For selective monitoring of hydrated electron reaction with lindane, experiments were performed in a N2-saturated 60 M i-PrOH solution that scavenges •OH radicals and •H atoms to form the relatively inert radical (reactions (3.15) and (3.16)). The rate of decay of the transient absorbance at 700 nm was used to monitor the reaction of hydrated electrons with lindane.

− Lindane + eaq → Product (3.36)

The decay curves fit an exponential rate law, indicating first order kinetics behavior. Plotting these pseudo-first order values against lindane concentration, a second- order rate constant of k = (1.26 ± 0.04) × 1010 M˗1 s˗1 was obtained. Thus, the rate constant for hydrated electrons reaction with lindane is of the order of diffusion- controlled reaction. This rate constant is approximately two order of magnitude higher than the reported rate constant of hydroxyl radical with lindane i.e. 2.9 ×108M−1 s−1, as reported by Nienow and co-workers (2008) and 7.5 × 108 M−1 s−1, as reported by Haag and Yao (1992). This rate constant is in good agreement with the value reported for reactions of hydrated electron with other chlorinated hydrocarbons; where values 1.4 ×

10 10 ˗1 ˗1 10 and 1.3 × 10 M s have been reported for CHCl3 and CCl4, respectively (Schmidt et al., 1995, Buxton et al., 1988).

100

3.3.9 Competition kinetics–Hydroxyl radical rate constant (•OH + Lindane)

Unlike hydrated electron, hydroxyl radical has small absorption in the accessible region that is not useful for determination of rate constant. Therefore, a competitive kinetics method was adopted to determine the absolute (second-order) rate constant of the hydroxyl radical (•OH) with lindane. The 2-chlorophenol (2-CP) which has well known second-order rate constant with •OH radical was used as reference compound. The second-order rate constant of •OH with lindane was determined by using equation (3.38)

(Haag and Yao, 1992);

k lindane lnlindane  lindane  t 0 (3.37) k2 CP ln 2  CP 2  CP    t  0

lnlindane  lindane k lindane t 0  k2  CP (3.38) ln 2CP 2 CP  t  0 where k(lindane) and k(2-CP) are the second-order rate constants for the reaction of •OH with lindane and 2-chlorophenol, respectively. [lindane]0 and [2-CP]0 represent the initial concentrations of lindane and 2-chlorophenol, respectively and [lindane]t and [2-

CP]t represent the concentrations of lindane and 2-CP, respectively, at reaction time “t”.

From the above reaction (3.38), the second-order rate constant of lindane with •OH was determined to be 6.8 × 108 M−1 s−1, which is consistent with the literature values of 7.5 ×

108 M−1 s−1 (Haag and Yao, 1992) and 2.9 ×108M−1 s−1 (Nienow et al., 2008), where the

• OH radical was generated by photo-Fenton’s or UV/H2O2 reactions, respectively.

Comparing these rate constant data, it appears that radiolytic decay of lindane occurs mainly by reductive pathway via hydrated electron reaction.

101

3.3.10 Variation of solution pH during irradiation

Figure 3.10 describes pH values of lindane solution before and after gamma radiation. It is evident that solution pH decreases during gamma irradiation. Higher absorbed dose resulted in more distinct decrease in pH values. The decrease of solution pH values is probably due to the formation of organic acids during the degradation

+ process (Wren and Glowa, 2000). Similarly, the hydronium ion (H3O ) produced in the irradiation process (equation (1.6)) may also cause the reduction of pH. Figure 3.10 shows that at absorbed dose of 2000 Gy, the pH of the solution was reduced from 6.8 to

4.8.

102

7.0

6.5

6.0

5.5

5.0

4.5

4.0 0 500 1000 1500 2000 Absorbed dose (Gy)

Fig. 3.10. Variation of pH value with the increasing radiation dose in gamma radiolytic

lindane decay. Experimental conditions: [lindane]0 = 3.43 µM, N2-saturated.

103

3.3.11 Identification of by-products and possible reaction pathways

Gas chromatography-mass spectrometry (GC/MS) analysis of the irradiated lindane solutions showed a number of reductive as well as oxidative intermediates. The major organic intermediates identified during gamma irradiation included pentachlorocyclohexene (PeCCH), tetrachlorocyclohexene (TeCCH), 1,4- cyclohexanedione and various short chain organic acids, such as formic, acetic, succinic and tartaric acids. Among these by-products, pentachlorocyclohexene (PeCCH) and tetrachlorocyclohexene (TeCCH) are the widely reported intermediates formed when lindane is decomposed in various processes, such as photocatalysis, microwave decomposition and reduction by zero valent iron (Salvador et al., 2002, Wang et al.,

2009; Senthilnathan and Philip., 2010). Due to the presence of high electronegativity chlorine atoms in lindane, it is easily susceptible to attack by the hydrated electrons

(Taghipour and Evans, 1997). The hydrated electrons of water radiolysis can remove one or more electrons from lindane in single or multiple steps, resulting in the generation of different kinds of reaction intermediates. The PeCCH can simply be obtained by eliminating one HCl molecule (dehydrochlorination reaction) from lindane by the reaction of hydrated electron. The PeCCH molecule on further dehydrochlorination may lead to the formation of TeCCH (Li et al., 2011). According to Wang and co-workers

(Wang et al., 2009), lindane may be reduced by the electron, forming a double bond by eliminating one HCl (dehydrochlorination) to give PeCCH or it may lose two Cl−

(dichloroelimination) to form tetrachlorocyclohexene (TeCCH). Similar intermediates have been reported during photolysis of lindane in the presence of polyoxometalate

3− − • (PW12O40 ), where aqueous electron (e aq) and hydroxyl radicals ( OH) both exist in the

104

reaction mixture (Antonaraki et al., 2010). Dechlorination reactions may eventually introduce double bond and unsaturation in lindane molecule and thus oxidation reactions by •OH radicals are involved. Hydroxyl radicals generally have high rate constant for reactions with aromatic compounds and unsaturated alkenes (Güsten et al., 1981). •OH radical has strong electrophilic nature and it may abstract hydrogen atom from chlorinated hydrocarbons (Spinks and Woods, 1990). Similarly, lindane may also be decomposed by abstraction of hydrogen atom by the •OH. However, lindane is less prone to oxidation due to the non-aromatic and fully saturated structure (Dionysiou et al.,

2000). The dehydrochlorination of lindane by the hydrated electron may create aromaticity in the molecule. The further degradation of aromatic compound leads to ring cleavage reaction generating a mixture of low-molecular-weight organic acids

(Brinkmann et al., 2003; Fan et al., 2011). The identification of different oxidative products, such as 1,4-cyclohexanedione and the organic acids suggests that oxidative pathway is also involved in the gamma radiolytic degradation of lindane. The proposed degradation pathway of lindane is given in scheme 1.

105

Scheme 1. Proposed lindane degradation pathway by gamma-ray irradiation.

106

3.3.12 Removal efficiency of lindane

Removal efficiency of pollutant can be expressed in several different ways. The simplest way is the percent (%) removal of pollutant which is obtained by comparing the concentration of pollutant before and after irradiation as given in equation (3.39):

C -C Percent removal efficiency  = 0 × 100 % (3.39) Co where C0 is the initial concentration and C is the concentration after radiation treatment.

Another common approach used for determining the process efficiency is the G value. i.e., the number of solute molecules decomposed per 100 eV of radiation energy absorbed. The following equation (reaction 3.40) can be used to calculate G value at various absorbed doses (Lewins et al., 1991; Basfar et al, 2005a):

ΔR N  Radiation chemical yield G, species/100 eV = A (3.40) D 6.24 × 1016  where ΔR is the change in lindane concentration (mol/L) at the given absorbed dose, NA is Avogadro’s number (6.02 × 1023 molecules/mol), D is radiation dose in Gy and 6.24 ×

1016 is conversion factor from Gy to 100 eV/L.

The G-values obtained under our experimental conditions are presented in Tables

3.7a - 3.7c. G-values provide some useful information about the nature of the radiolytic reactions. For example, for any radiolytic decay, the G-values less than 0.31 indicate that there is no radical chain reaction occurring in the system (Getoff and Lutz, 1985). Thus, lower G-values (˂ 0.31) observed under our experimental conditions (Tables 3.7a - 3.7c) shows that no radical chain reactions are involved in the present study. From the results given the Tables 3.7a - 3.7c, it can be seen that the G-values decreased with increase in

107

accumulated absorbed dose. The reduction in G-value with increasing absorbed dose is due to competition of lindane molecules with the intermediates resulted from lindane for reactive radicals. The concentration of intermediates is increased with increasing absorbed dose whereas the concentration of lindane is decreased. Therefore, at accumulated absorbed dose, the possibility of reactive radicals to react with intermediates molecules rather than with parent compound increases and hence the G-values decreases.

This trend has been observed in the radiolytic decay of several other compounds (Lin et al., 1995; Basfar et al., 2005a; Yu et al., 2008).

Still another quantitative way to represent efficiency of solute removal is to calculate dose constants (or decay constant) ‘k’ that explain the concentration of solute removed per unit dose. Under our experimental conditions, the concentration of reactive radicals is in excess to organic solute which can be expressed by pseudo-first order kinetics. Under pseudo-first order kinetics, the rate of decomposition is directly proportional to the amount of undecomposed material as mathematically given in equation 3.11 (in section 3.5). Equation 3.11 is generally used to calculate the dose constants by determining the slope of the plot of natural logarithm of the solute concentration (ln[solute]) vs. absorbed dose (Gy). These dose constants can be used to calculate the dose required for some specific percentage reduction in solute, e.g. for 50 and 90% reduction in the solute concentration as mathematically expressed in reactions

3.41 and 3.42, respectively (Basfar et al, 2005b).

D0.5 = ln (2)/k (3.41)

D0.9 = ln (10)/k (3.42)

108

The values of dose constants ‘k’ obtained under our experimental conditions are provided in Tables 3.8a - 3.8e. From the data shown in Tables 3.8a - 3.8e, it can be observed that the dose constants ‘k’ change with the change of the rate of lindane degradation under different experimental conditions. The dose constant (k) is directly associated with the rate of solute degradation and inversely related to D0.5 and D0.9. From the results given in tables 3.8a-3.8e it can be seen that higher the rate of lindane degradation under particular experimental conditions, higher the values of k and lower the values of D0.5 and D0.9 and vice versa.

109

Table 3.7a. G-values for lindane removal in the presence of various organic and inorganic substances. Experimental − − − − 2− conditions: [lindane]0 = 3.43 µM, [NO2 ] = [NO3 ] = [Cl ] = [HCO3 ] = [CO3 ] = 1 mM, N2-purged for 20 min.

Absorbed G-value

dose (Gy) − − − − 2− Blank NO2 NO3 Cl HCO3 CO3 CHCl3 Humic acid 2-CP

(3.43 µM) (5mg/L) (3.43 µM)

250 0.0699 0.0117 0.0167 0.0619 0.0753 0.0787 0.0502 0.0588 0.0586

500 0.0544 0.0133 0.0167 0.0477 0.0569 0.0586 0.0460 0.0502 0.0502

1000 0.0347 0.0146 0.0159 0.0330 0.0372 0.0376 0.0309 0.0326 0.0343

1500 0.0262 0.0125 0.0139 0.0245 0.0270 0.0273 0.0245 0.0251 0.0260

2000 0.0203 0.0109 0.0121 0.0192 0.0209 0.0209 0.0194 0.0199 0.0201

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Table 3.7b. G-values for lindane removal in N2O-saturated, aerated and N2-saturated solutions as well as in N2-saturated

solutions containing H2O2, t-BuOH and i-PrOH. [lindane]0 = 3.43 µM. Absorbed G-value

dose N2 N2O Aerated H2O2 H2O2 H2O2 t-BuOH i-PrOH (Gy) (5 mM) (10 mM) (20 mM) (60 mM) (60 mM)

250 0.0699 0.0101 0.0112 0.0355 0.0167 0.0055 0.0468 0.0385

500 0.0544 0.0083 0.0134 0.251 0.0134 0.0075 0.0418 0.0376

1000 0.0347 0.0083 0.0127 0.0167 0.0104 0.0075 0.0326 0.0284

1500 0.0262 0.0083 0.0111 0.0139 0.0097 0.0069 0.0245 0.0231

2000 0.0203 0.0074 0.0102 0.0125 0.0061 0.0059 0.0192 0.0182

111

Table 3.7c. G-values for lindane removal in gamma rays irradiated N2-saturated aqueous solution; Effect of initial lindane concentration. pH = 6.8. Absorbed dose G-value

(Gy) [lindane]0 = [lindane]0 = [lindane]0 = [lindane]0 =

0.343 µM 0.694 µM 1.715 µM 3.434 µM

250 0.1004 0.0971 0.0753 0.0669

500 0.0669 0.0653 0.0602 0.0544

1000 0.0397 0.0393 0.0381 0.0347

1500 0.0276 0.0274 0.0270 0.0262

2000 0.0208 0.0208 0.0206 0.0203

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Table 3.8a. Observed dose constant (k), D0.5 and D0.9 for lindane removal in N2 saturated gamma ray irradiated aqueous solutions containing 1 mM of different inorganic anions.

− − − − 2− Parameters Blank NO2 NO3 Cl HCO3 CO3

k (Gy−1) 1.76 × 10−3 3.87× 10−4 4.44 × 10˗4 1.27 × 10−3 2.31 × 10−3 2.55 × 10−3

2 3 3 2 2 2 D0.5 (Gy) 3.95 × 10 1.79 × 10 1.56 × 10 5.48 × 10 3.01 × 10 2.72 × 10

3 3 3 3 2 2 D0.9 (Gy) 1.31 × 10 5.96 × 10 5.18 × 10 1.82 × 10 9.99 × 10 9.03 × 10

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Table 3.8b. Observed dose constant (k), D0.5 and D0.9 for lindane removal in N2O-saturated, aerated and in N2-saturated

containing CHCl3, humic acid and 2-CP solution.

N2 N2O Air CHCl3 Humic acid 2-CP

Parameters (3.43 µM) (5mg/L) (3.43 µM)

k (Gy−1) 1.76 × 10−3 2.24 × 10−4 3.40 × 10−4 1.33 × 10−3 1.48 × 10−3 1.66 × 10−3

2 3 3 2 2 2 D0.5 (Gy) 3.90 × 10 3.10 × 10 2.04 × 10 5.19 × 10 4.69 × 10 4.17 × 10

3 4 3 3 3 3 D0.9 (Gy) 1.31 × 10 1.03 × 10 6.77 × 10 1.72 × 10 1.56 × 10 1.39 × 10

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Table 3.8c. Observed dose constant (k), D0.5 and D0.9 for gamma ray lindane removal in aqueous solutions containing

inorganic oxidant (H2O2) and common radical scavengers (t-BuOH and i-PrOH).

Parameters H2O2 t-BuOH i-PrOH

(5 mM) (10 mM) (20 mM) (60 mM) (60 mM)

k (Gy−1) 4.23 × 10−4 2.85 × 10−4 1.73 × 10−4 1.29 × 10−3 1.31 × 10−3

3 3 3 2 2 D0.5 (Gy) 1.64 × 10 2.43 × 10 4.01 × 10 5.36 × 10 5.29 × 10

3 3 4 3 3 D0.9 (Gy) 5.45 × 10 8.07 × 10 1.33 × 10 1.78 × 10 1.76 × 10

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Table 3.8d. Observed dose constant (k), D0.5 and D0.9 for gamma ray lindane removal at four different initial solute concentrations (0.343, 0.694, 1.715 and 3.43 µM). Experimental conditions: pH = 6.8. Parameters Lindane

0.343 µM 0.694 µM 1.715 µM 3.434 µM

k (Gy−1) 2.90 × 10−3 2.64 × 10−3 2.20 × 10−3 1.76 × 10−3

2 2 2 2 D0.5 (Gy) 2.39 × 10 2.63 × 10 3.15 × 10 3.95 × 10

2 2 3 3 D0.9 (Gy) 7.94 × 10 8.73 × 10 1.05 × 10 1.31 × 10

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Table 3.8e. Observed dose constant (k), D0.5 and D0.9 for gamma ray lindane removal at three different pH (4, 6.8 and 8).

Experimental conditions: [lindane]0 = 3.434 µM. Parameters pH = 4 pH = 8 pH = 6.8

k (Gy−1) 3.53 × 10−4 1.07 × 10−3 1.76 × 10−3

3 2 2 D0.5 (Gy) 1.97 × 10 6.46 × 10 3.95 × 10

3 3 3 D0.9 (Gy) 6.53 × 10 2.15 × 10 1.31 × 10

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3.4 Degradation of Lindane by Photochemical Oxidation

− In this section, degradation and kinetics of peroxymonosulfate (PMS, HSO5 ) based photochemical oxidation of lindane was investigated at lab-scale experiments. The experiments were carried out under the following conditions:

 Degradation of lindane by solely PMS or ferrous iron (Fe2+) and direct UV

photolysis

 PMS activated by ferrous iron: Fe2+/PMS system

 Fe2+/PMS system assisted by tube-light radiation: Tube-light/Fe2+/PMS system

 PMS activated by UV radiation: UV/PMS system

 UV/PMS system assisted by Fe2+: UV/Fe2+/PMS system

3.4.1 Degradation of lindane by solely PMS or Fe2+ and direct UV photolysis

Initially, control experiments were conducted to determine the lindane removal by solely PMS or ferrous iron (Fe2+) or direct UV photolysis and the results are presented in

Fig. 3.11a. The results showed that PMS or Fe2+ alone as well as direct UV photolysis did not give significant degradation of lindane in 3 hour reaction time.

− PMS (HSO5 ) is a strong oxidizer with standard oxidation–reduction potential

(E0) of 1.82 V and it is capable of oxidizing many organic compounds under various activation conditions. However, PMS is a stable reagent and reaction rates with organic molecules are usually very slow such that no appreciable degradation of lindane was observed in the present study in three hour when 300 µM of PMS was used.

Ferrous iron (Fe2+) does not show any significant degradation of pollutant when it is applied to the system, individually. The Figure 3.11a shows that at the initial Fe2+ concentration of 300 µM, only 6% lindane was removed in three hours reaction time. It

118

has been reported that less than 5% pollutant removal can be achieved when sufficient amount of Fe2+ is present for the degradation of PCBs, propachlor (Liu et al., 2012) and xanthene dye (Wang and Chu, 2011) in aqueous solution.

The result given in Figure 3.11a shows that lindane is resistant to direct UV photolysis and the decrease in lindane concentration is very small at reasonably high UV dose. For exposure time of three hours UV radiation, only 8% lindane was removed from

3.43 μM solution of lindane. It is well known that a molecule containing chromophore or double bond can easily absorb light energy in the UV range (Dantas et al., 2010).

However, lindane molecule with fully saturated structure cannot absorb UV radiation directly and hence no significant decomposition can be observed in direct UV photolysis.

This is in agreement with the UV absorption spectrum of lindane shown in Figure 3.11b, indicating no UV absorption in this region. Previous studies have shown comparable results for direct photolysis of other chlorinated compounds. Zalazar and co-workers

(2007) found that dichloroacetic acid gave no sign of degradation under direct UV photolysis. Li and co-workers (2012) also found that no observable decrease in the concentration of monochloro acetic acid occurred upon direct photolysis. The small amount of lindane degraded in the direct UV photolysis can be attributed to the production of some reactive species resulting from photolysis of water molecules

(Sheoran, 2008).

Accordingly, it can be concluded that lindane degradation by direct UV photolysis, solely PMS or ferrous ion is negligible and thus all these systems cannot be applied for decontamination of lindane contaminated waters, individually.

119

(a)

1.0

PMS 0.8 2+ Fe UV photolysis

(lindane)

0 0.6

C/C

0.4

0 20 40 60 80 100 120 140 160 180 200 Reaction time (min)

(b)

Fig. 3.11a. Degradation of lindane by solely PMS, ferrous iron (Fe2+) or direct UV

2+ photolysis. Experimental conditions: [Lindane]0 = 3.43 μM, [PMS]0 = [Fe ]̥0 =

300 μM, (b) UV Absorbence spectra of lindane solution at C0 = 3.43 μM.

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3.4.2 PMS activated by Fe2+: Fe2+/PMS system

Preliminary results showed that neither Fe2+ nor PMS alone lead to any significant degradation of lindane at ambient conditions in three hours. Both Fe2+ and PMS, being fairly stable reagents can react with lindane very slowly when present in solution, individually. Many transition metals, especially divalent metals (M2+) may act as electron donors to catalyze the decomposition of PMS through one-electron transfer reaction analogous to the Fenton initiation reaction. Ferrous ion (Fe2+) has many advantages as an oxidant activator because it is cheap, less toxic and abundant metal in nature (Nfodzo and

•− − Choi, 2011). The formation of sulfate radical (SO4 ) from PMS (HSO5 ) in this Fenton- like process (Fe2+/PMS system) has been recently explored (Anipsitakis and Dionysiou,

•− 2003, 2004b). The major reactions involved in the production and consumption of SO4 using Fe2+/PMS system is given in reactions 3.43-3.46;

2+ − 3+ •− − 4 −1 −1 Fe + HSO5 → Fe + SO4 + OH (k = 3.0 × 10 M s ) (3.43)

2+ •− 3+ 2− 8 −1 −1 Fe + SO4 → Fe + SO4 (k = 3.0 × 10 M s ) (3.44)

•− − •− 2− + 5 −1 −1 SO4 + HSO5 → SO5 + SO4 + H , (k < 1.0 × 10 M s ) (3.45)

•− •− − 8 −1 −1 SO4 + SO4 → S2O82 , (k = 4 × 10 M s ) (3.46)

Although Fe2+ is an activator of PMS, the reaction is kinetically slow and usually

•− it takes long time for completion. Large quantities of SO4 are lost due to its scavenging

2+ − •− reactions with Fe , PMS (HSO5 ) or other SO4 as shown in reactions (3.44), (3.45) and

(3.46), respectively. The ferric iron (Fe3+) produced in reactions (3.43) and (3.44) cannot activate PMS and thus, additional step is required for regeneration of the Fe2+ catalyst.

•− 2+ Thus reaction (3.43) is a limiting step towards the production of SO4 in the Fe /PMS

(dark) system. As a result, only 11% lindane removal can be achieved in three hour, using

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300 μM initial concentration of Fe2+ and PMS, each. The extent of lindane degradation by Fe2+/PMS system, however, increased to 25% when prolong reaction time (20 h) is given to the reaction (results not shown). The results of lindane removal by Fe2+/PMS system in three hour reaction time are shown in Figure 3.12a. It was found that the degradation efficiency of lindane increased by increasing the initial concentration of both

Fe2+ and PMS up to certain extent until the molar ratios of lindane: ferrous iron: oxidant

(LFO) is 1:87:87, respectively. Starting with the initial lindane concentration of 3.43 μM,

300 μM Fe2+ and 300 μM PMS was chosen as optimum doses for removal of lindane in aqueous solution. The results showed that 1:1 catalyst/oxidant molar ratio is the optimum value and deviation from the 1:1 molar ratio led to a decrease in the removal efficiency of the Fe2+/PMS system. The 1:1 catalyst/oxidant molar ratio was found as optimized value in other studies also in the literature (Gupta and Sutar, 2008; Rastogi et al., 2009). As

2+ − •− indicated in reactions 3.52 and 3.53, both Fe and PMS (HSO5 ) can scavenge SO4 when these reagents are present in excess amount. Thus moderate quantity of both Fe2+ and PMS are needed for the optimum removal of lindane in the Fe2+/PMS system.

2+ Fe is recognized as a good activator of H2O2 but it exhibits only limited ability in the activation of PMS for the decontamination of pollutants. Several researchers found that cobalt (Co2+) catalyst is an efficient activator of PMS, however, the use of Co2+ has severe environmental concerns and thus Co2+ cannot be a good alternative in practical applications (Chan et al., 2009; Do et al., 2009). Therefore, there is a need to explore new methods that can improve the efficiency of the Fe2+/PMS system, which is comparatively environmental friendly method for the removal of pollutants. Several attempts were made for enhancing the efficiency of Fe2+/PMS process and one such attempt is the application

122

of light energy that may enhance the catalytic efficiency of transition metal ions in the oxidation of pollutants. Two typical energy sources i.e. UV-C light and visible light

(ordinary room tube-light) were employed for activation of Fe2+ catalyst and the results are discussed in the next section.

3.4.3 Fe2+/PMS system assisted by tube-light radiation: Tube-light/Fe2+/PMS system

Very interesting results were obtained when tube-light assisted Fe2+/PMS system was employed for the degradation of lindane in water. In the presence of tube light radiation, 3.43 µM lindane solution containing 300 μM Fe2+ and 300 μM PMS at pH 3.0, led to 40% overall lindane removal in three hours reaction time. Literature study shows that iron-mediated (photo-Fenton) process is sensitive to light up to wavelengths ≤ 600 nm (Malato et al., 2002) and hence the efficiency of the Fe2+/PMS system is increased in the presence of tube-light radiation. The enhanced removal efficiency of the tube- light/Fe2+/PMS system can be attributed to the regeneration of iron catalyst (Fe2+) (Hislop and Bolton, 1999) resulted from photochemical effects of the tube-light radiation. The visible light (including tube-light radiation) consists of a wide range of wave-lengths (λ =

300-500 nm) and a portion of this energy may be used for the reduction of Fe3+ into Fe2+ ions (Pignatello, 1992; Zepp et al., 1992; Chong et al., 2010). The reversion cycle of

2+ 3+ 2+ •− Fe (aq) → Fe (aq) → Fe (aq) continuously generates SO4 , provided that the concentration of PMS in the system is substantial. Alternatively, the absorption of energy

− •− photons by HSO5 ions may directly result in the generation of SO4 (reaction 3.47) or it

− 3+ 2+ may excite HSO5 ion which can be used for the reduction of Fe to Fe ion

(Chevaldonnet et al., 1986).

123

The regeneration of the Fe2+ (aq) from Fe3+ (aq) is the rate-limiting step in the catalytic iron cycle, if small amount of iron is present in the solution (Chong et al., 2010).

In the present study, varying concentrations of Fe2+ and PMS were employed and the optimum concentrations of Fe2+ and PMS in the tube-light/Fe2+/PMS system were selected. The results of lindane removal by tube-light/Fe2+/PMS system are shown in

Figure 3.12a. The results showed that the degradation efficiency of lindane increased by increasing the initial concentration of both Fe2+ and PMS up to certain extent which ultimately diminished after some optimum values. In Figure 3.12b, variation of rate constant ‘k’ with the increasing concentration of Fe2+ and PMS for the Tube- light/Fe2+/PMS system is shown. At the initial lindane concentration of 3.43 μM, the optimum molar ratio of lindane: ferrous iron: oxidant (LFO) is found to be 1:87:87, respectively.

124

(a)

1.0

0.8

0.6

(lindane)

C/C 0.4

Fe2+/PMS 2+ 0.2 Tubelight/Fe /PMS

0.0 0 20 40 60 80 100 120 140 160 180 200 (b) Reaction time (min)

0.16

0.14

0.12

)

-1

(h 0.10 k [Fe2+] [PMS] 0.08

0.06

Rate constant, Rate 0.04

0.02

0.00 0.0 0.5 1.0 1.5 2.0 2.5

2+ Concentration of PMS or Fe (mM)

Figure 3.12. (a) Lindane degradation as a function of reaction time in the Dark/Fe2+/PMS

2+ and Tube-light/Fe /PMS systems. Experimental conditions: [lindane]̥0 = 3.43 2+ μM, [PMS]̥0 = 300 μM, [Fe ]= 300μM and pH = 3.0. (b) The variation of rate constant ‘k’ (h−1) with the increasing concentration of Fe2+ and PMS for the Tube- light/Fe2+/PMS system is shown.

125

The results given in Figure 3.12b showed that the molar ratios of Fe2+ and PMS played a significant role in lindane degradation in the tube-light/Fe2+/PMS system. The results showed that molar ratio of the catalyst/oxidant is a crucial factor in the tube- light/Fe2+/PMS system. It was found that the system give maximum efficiency when the catalyst/oxidant molar ratios is 1:1. Variation from 1:1 molar ratio led to a decrease in the lindane removal efficiency in the tube-light/Fe2+/PMS system. Our results of the 1/1 catalyst/oxidant molar ratios is consistent with the findings of other researchers (Wang and Chu, 2011). Rastogi and co-workers (2009) tested the efficiency of the Fe+2/PMS system for PCBs degradation in aqueous and sediment systems and 1:1 was found to be the optimum ratio among the oxidant and catalyst.

3.4.4 PMS activated by UV radiation: UV/PMS system

UV radiation is generally considered as an efficient activator of oxidants that are employed in water treatment for the degradation of pollutants (Hijnen et al., 2006; Li et al., 2012). In the present study, UV/PMS system was applied for the removal of lindane and the results are shown in Figure 3.13. The results showed that although direct photolysis did not result in significant lindane degradation, the addition of PMS resulted in rapid degradation of lindane.

•− Highly reactive sulfate free radicals (SO4 , E˚ = 2.5-3.1 V) can be generated by

− activation of PMS (HSO5 ) with UV radiation via reaction 3.47 (Anipsitakis and

•− Dionysiou, 2004b; Shukla et al., 2010). The SO4 generated can act as strong oxidant for organic compounds via hydrogen abstraction, addition to unsaturated carbon or by electron removal process (Neta and Zemel, 1977; Huie et al., 1991).

126

Different initial concentrations of PMS were used for removal of lindane in the

UV/PMS system and the corresponding degradation results are shown in Fig. 3.14a. The pseudo-first order rate constants ‘k’ obtained at using different initial concentrations of

PMS in the UV/PMS process are shown in Table 3.9 and Figure 3.14b. The results given in Figure 3.14a showed that for a given concentration of lindane, the rate of lindane degradation increased with increasing the initial concentration of PMS. Such an increase in the degradation rate with increasing initial PMS concentration was observed in

UV/PMS oxidation of several other pollutants (Chen et al., 2007; Antoniou et al., 2010a).

The enhanced degradation efficiency of the UV/PMS process with increasing initial PMS

•− concentration is due to increased number of SO4 radicals produced when higher

•− concentration of PMS is used. Apart from reaction with organic pollutants, the SO4 may

− react with PMS (HSO5 ) or it may also react with itself to generate unreactive species

(reactions 3.44 and 3.45). However, the scavenging effect of these reactions is comparatively small at lower PMS concentration while it increases with the increase of

PMS concentration. This is in accordance with the results given in Figure 3.14a, which shows that the relative increase in the degradation rate with increasing PMS concentration is high in the beginning, while it gradually slows down at higher PMS concentrations. Under the optimum conditions, using 300 μM PMS, 97% lindane was removed in 3 h of photolysis, when the lindane: PMS molar ratio is 1:87. The various

•− reactions involved in the production and consumption of SO4 by the UV/PMS system is given in the following reactions (Anipsitakis and Dionysiou, 2004b);

•− Generation of SO4 by UV radiations:

− •− • HSO5 + hν → SO4 + OH (3.47)

127

•− Scavenging reactions of the SO4 , besides the reactions 3.45 and 3.46 include;

•− + 2− • SO4 + H2O → H + SO4 + OH (3.48)

(at all pH values)

•− − 2− • 7 −1 −1 SO4 + OH → SO4 + OH (k= 1.4-7.3 ×10 M s ) (3.49)

(mostly in alkaline pH)

• − •− OH + HSO5 → SO5 + H2O (3.50)

128

1.0

0.8 Direct UV-Photolysis UV/PMS

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Figure 3.13. Lindane degradation as a function of time of photolysis in the (A) direct UV

photolysis and (B) UV/PMS processes. Experimental conditions: [lindane]̥0 = 3.43 μM, [PMS] = 300 μM and pH = 6.0.

Table 3.9. Pseudo-first order rate constants for UV/PMS photochemical degradation of lindane (3.43 µM) in the presence of various amounts of PMS.

−1 2 No. [PMS]0 (µM) kap (min ) R

1 100 4.78 × 10−3 0.998

2 300 1.32 × 10−2 0.978

3 500 2.53 × 10−2 0.966

4 1000 3.40 × 10−2 0.984

5 2000 4.17 × 10−2 0.988

129

(a)

[PMS]=0.1 mM 1.0 [PMS]=0.2 mM [PMS]=0.5 mM [PMS]=1.0 mM 0.8 [PMS]=2.0 mM

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

(b)

0.05

0.04

)

-1

(min 0.03

0.02

Rate constant, constant, Rate 0.01

0.00 0.0 0.5 1.0 1.5 2.0 2.5 Concentration of PMS (M)

Figure 3.14. (a) Effect of initial concentration of PMS on the UV/PMS Photochemical

degradation of lindane. Experimental conditions: [lindane]̥o = 3.43 μM, pH = 6.0. (b) Variation of rate constant ‘k’ (min−1) with the increasing PMS concentration is shown. 130

3.4.5 UV/Fe2+/PMS system

As stated earlier, both UV radiation and ferrous iron (Fe2+) can activate PMS

•− independently to generate highly reactive SO4 . In addition, UV radiation may also be used for regeneration of the Fe2+ from Fe3+ ions that consequently activate PMS and thus

•− further increase the concentration of SO4 (Zepp et al., 1992; Bossmann et al., 1998).

This system was applied for degradation of lindane as well. The result of UV/Fe2+/PMS system for the removal of lindane is shown in Figure 3.15. The results showed that the removal efficiency of the Fe2+/UV/PMS system increased with the increasing concentration of both PMS and Fe2+ up to certain limit until the lindane/catalyst/oxidant molar ratio is 1: 87: 15, respectively. However, above this optimum level, only slight increase in the lindane removal efficiency occurred with further increase in the concentration of PMS. The variation of the observed degradation rate constants, k with the increasing concentration of both PMS and Fe2+ is shown in Figure 3.16a and Figure

3.16b. When the concentrations of PMS and Fe2+ exceed their optimum values, the rate

•− of production of SO4 radicals and its consumption by the scavenging reactions become

•− comparable and hence only slight increase in concentration of SO4 is observed

(Anipsitakis and Dionysiou, 2003; Liang et al, 2004; Fernandez et al., 2004).

Surprisingly, a slight decrease in the lindane removal efficiency can be observed when the amount of Fe2+ surpasses the optimum concentration limits. At such reaction stage,

•− the scavenging reactions surmount the production rate of SO4 and hence a slight decline in the lindane removal efficiency can be observed. Such a decrease in the degradation rate of pollutants with the increasing Fe2+ concentration is supported by the results of other researchers (Tamimi et al., 2008; Pagano et al., 2011). In the present study, the

131

optimum values for the oxidant/catalyst concentrations were attained at 6:1 molar ratio, which gave 97% lindane removal in 40 min.

It is evident from the results that significantly enhanced degradation results were achieved when UV radiation and ferrous iron (Fe2+) were simultaneously involved in the activation of PMS for the removal of lindane. Compared to the individual systems of

Fe2+/PMS and UV/PMS which required 3655 and 91 min, respectively for 90% lindane removal, the combined UV/Fe2+/PMS system requires only 24 mints for the removal of same quantity of lindane from aqueous solution. The conversion of Fe2+ into Fe3+ ions

•− 2+ (reaction 3.51) may limit the production of SO4 in the Fe /PMS system. The reversion

3+ 2+ •− of Fe into Fe can sustain the generation of SO4 radicals, provided that the concentration of PMS in the system is substantial. In the UV/Fe2+/PMS system, the UV radiation is used in the regeneration of active Fe2+ catalyst from the Fe3+ as shown in reaction 3.51 (Hislop and Bolton 1999). Thus the synergistic effect of the UV radiation and ferrous iron (Fe2+) can be the plausible explanation for the greatly enhanced pollutant removal efficiency of the UV/Fe2+/PMS process.

Fe(OH)2+ +hν→ •OH + Fe2+ (3.51)

132

1.0

0.8 Direct UV-Photolysis UV/Fe2+/PMS

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min) Figure 3.15. Lindane degradation as a function of time of photolysis in the (A) direct UV 2+ photolysis and (B) Fe /UV/PMS processes. Experimental conditions: [lindane]̥0 = 3.43 μM, [PMS] = 300 μM, [Fe2+] = 50 μM and pH = 3.0.

133

(a)

3.0

2.5

)

-1 2.0

1.5

1.0

0.5

Rate constant, (min k

0.0

0.0 0.2 0.4 0.6 0.8 1.0 1.2

2+ Concentration of Fe (mM)

(b)

0.18

0.16

0.14

) -1 0.12

0.10

0.08

0.06

0.04

Rate constant, k (min constant, Rate

0.02

0.00

0.0 0.2 0.4 0.6 0.8 1.0 1.2 Concentration of PMS (mM) Figure 3.16. (a) Variation of rate constant ‘k’ (min−1) with the increasing concentration of 2+ 2+ Fe for UV/Fe /PMS system at [PMS]0 = 300 μM. (b) Variation of rate constant ‘k’ (min−1) with increasing concentration of PMS for UV/Fe2+/PMS system at 2+ [Fe ]0 = 50 μM. Experimental conditions: [lindane]̥0 = 3.43 μM, pH = 3.0.

134

3.4.6 Kinetics of UV/PMS oxidation

UV/PMS system was selected as a model sulfate radical based advanced oxidation process (SRB-AOP) for the decomposition of lindane in aqueous solution. For kinetics studies, the experiments were conducted at various initial concentrations of lindane and at different pH values.

The mechanisms of oxidation by UV/PMS have been investigated extensively, and it has been found that the rate of degradation of an organic compound using UV/PMS process results from the contribution of two pathways: direct photolysis and the reactive radicals attack (Antoniou et al., 2010b; Yang et al., 2010):

dC   k C  k OH  C  k  SO   C UV   OH C   SO  C 4   dt 4 (3.52)

−• • •− where k•OH-C and kSO4 -C are the second-order rate constants for OH and SO4 reactions

• •− with lindane (C), respectively. Because OH and SO4 radical concentration can be

• −• assumed to be constant over the range of reaction, the products (k•OH-C [ OH]) and (kSO4 -

•− C [SO4 ]) are almost constant and can be considered pseudo-first order rate constants

−• kˈ•OH-C and kˈSO4 -C, respectively (Benitez et al., 2006). Thus, the following equation may be used to describe the degradation of lindane ‘C’ during UV/PMS process:

dC   kUV  C  k  C  k    C dt OH C SO4 C dC  (3.53)  kUV  k   k    C dt  OH C SO4 C  dC  kC  dt

Integrating between ‘0’ to ‘t’ with corresponding concentrations of ‘C0’ and ‘C’ yield;

135

C ln kt (3.54) C0  where ‘k’ represents the pseudo-first order rate constant for the overall degradation of the compound ‘C’ during the UV/PMS treatment. According to equation (3.54), a plot of

−ln[C]/[C]0 versus reaction time should lead to straight lines, with a slope equal to rate constant, k. From the previous discussion (also shown in Figure 3.11a), it is clear that direct UV photolysis does not degrade lindane significantly, so, the kUV can be neglected in above equation (3.54) and thus ‘k’ represents the combined second-order rate constants

• •− −• of OH and SO4 with lindane i.e. k•OH-C and kSO4 -C. The second-order rate constant of

•OH with lindane was determined in the previous section (i.e., 3.3.9) using competition kinetics, which was equal to 6.8 x 108 M−1 s−1. The second-order rate constant of lindane

•− 9 −1 −1 with SO4 was subsequently determined to be 1.3 × 10 M s , demonstrating a

•− • comparable reactivity of SO4 and OH with lindane. The reference compound used was meta-toluic acid (m-TA), which has a known second-order rate constant of 2.0 × 109 M−1

−1 •− s with SO4 (Neta and Zemel, 1977).

3.4.6.1 Effect of initial concentrations of lindane

The influence of the initial lindane concentration was also studied. Figure 3.17a illustrates the degradation efficiency of lindane at different initial concentrations. The observed degradation rate constant ‘k’ values obtained for different initial lindane concentration is given in Table 3.10 and Figure 3.17b. Variation of observed degradation rate constant, k with change in initial solute concentration is shown in Figure 3.17c. As expected for the pseudo-first order reaction, the degradation rate increased but the observed degradation rate constant ‘k’ decreased with increase of the initial concentration 136

of lindane. Literature studies show that decomposition of many organic pollutants at lower concentration ranges follow first-order kinetics with respect to pollutant concentration (Tang and Tassos, 1997; Elkanzi and Bee Kheng, 2000; Gao et al., 2012).

The increase in the degradation rate with increasing solute concentration is due to the

•− increased number of solute molecules reacting with the reactive radicals (SO4 ) per unit time. The decrease in the observed dose constant ‘k’ with increasing initial solute

•− concentration is due to the decreased SO4 /lindane ratio at the fixed concentration of

PMS. Enhanced formation of intermediates at higher initial lindane concentrations may absorb part of the UV light, therefore, lowering the observed degradation rate constants at higher pesticide concentrations (Chelme-Ayala et al., 2010). A similar degradation trend was observed in the degradation of Acid Orange 7 using SR-AOP (Chen et al., 2007).

137

(a)

1.0 [lindane]= 6.8 M [lindane]= 3.4 M [lindane]= 0.68 M

0.8

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

(b)

5 [lindan] = 0.68 M, R2 = 0.9994 [lindan] = 3.4 M, R2 = 0.9916 4 [lindan] = 6.8 M, R2 = 0.9652

)

- ln(C/C

0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

138

(c)

0.040

0.035

)

-1 0.030

(min k 0.025

0.020

0.015

Rate constant, constant, Rate

0.010

0.005 0 1 2 3 4 5 6 7 8 Concentration of lindane (M)

Figure 3.17. Kinetics of UV/PMS photodegradation of lindane at different initial solute concentrations. (a) Change in lindane concentration with time of UV photolysis,

(b) Plots of –ln(C/C0) Vs. time of UV photolysis, (c) Variation of observed degradation rate constant, k with change in initial solute concentration.

Experimental conditions: [PMS]̥0 = 300 μM, and pH = 6.0.

Table 3.10. Kinetics data of lindane photodegradation with UV/PMS processes at different initial solute concentration ______

−1 C0 (µM) kobs (min ) t1/2 (min) r (n = 3) 0.68 0.0380 18.2 0.978

3.4 0.0253 27.4 0.995

6.8 0.0099 69.5 0.985

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3.4.6.2 Effect of solution pH

The pH value of the solution is usually an important factor in the study of pollutant degradation as pH variation may cause changes in the concentration of reactive species. The pH of the most frequently studied Fenton and photo-Fenton systems is limited to around 3, while higher pH can result in conversion of the reactive ferrous iron

(Fe2+) into ferric (Fe3+) precipitate and thus hindering the overall rate of pollutant decomposition (Ghaly et al., 2001). However, degradation study in a wide pH range can be performed in those systems where no iron catalyst (Fe2+) is being used. PMS, being the component of an acidic salt, automatically lowers the pH of the solution to acidic range and normally small amount of acid is needed in case of Fenton-like processes using

PMS as oxidant. The results of the UV/PMS process conducted for the degradation of lindane at three different pH values (4.0, 6.0 and 8.0) are presented in Figure 3.18. As can be seen, the rate of lindane degradation is the highest at neutral pH, while it decreases both at lower as well as at higher pH. The results further showed the extent of retardation effect is more prominent at higher pH as compared to lower pH side. The lower lindane degradation efficiency at alkaline pH values can be attributed to the conversion of the

•− • more reactive SO4 species into OH radical which take place at higher pH values

(Huang et al., 2002; Romero et al., 2010; Chu et al., 2011). The results obtained in this study are in agreement with the literature findings where the rate of degradation of methyl tert-butyl ether (Huang et al., 2002), diphenylamine, acetic acid (Criquet and

Karpel Vel Leitner, 2011) and textile dye (Madhavan et al., 2006) decreased with increase of the solution pH in various SR-AOPs. Also, the results obtained are consistent with other studies where the rate of pollutant degradation decreased with the increase of

140

the solution pH in the Fenton and photo-Fenton like processes (Tamimi et al., 2008). A slight decrease in degradation efficiency at acidic pH is most likely due to the scavenging

• •− + effects of OH/SO4 radicals by the H ion under such conditions (Spinks and Woods,

1990; Huang et al., 2009). Similarly, the standard reduction potential of hydroxyl radical

(•OH) is strongly reduced at higher pH and this may be another factor leading to a decreased removal efficiency at higher pH (Buxton et al., 1988). In addition, the formation of carbon dioxide resulted from the degradation of organic pollutant (lindane) could lead to the accumulation of bicarbonate and carbonate ions under alkaline solutions, which might inhibit organic pollutant oxidation (Xu et al., 1989). PMS alone can also interact with bicarbonate species, most likely inducing the generation of

•− percarbonate ions (Anipsitakis et al., 2005). The concentration of SO4 may also be decreased by reaction with hydroxide ions at higher pH values (reaction 3.57), thus lowering the rate of lindane degradation (Criquet and Karpel Vel Leitner, 2011).

Previous studies by Huang and Huang (2009) have indicated that PMS based process exhibit better performances at neutral pH and this behavior can make the PMS oxidant as the best option for degradation of organic pollutants in real-world applications.

Our result is consistent with the results reported by Zheng and Richardson (1995) and is coincidentally similar to the findings of Huang and Huang (2009).

3.4.6.3 Effect of humic acid on the UV/PMS system

Humic acids (HAs) are yellow- to black-colored macromolecular substances that constitute a considerable fraction of natural organic matter (NOM) in surface waters.

NOMs can have an effect on the rate of photodegradation of organic pollutants by virtue of •OH radical scavenger. It may also absorb light (inner filter effect) and can undergo

141

different photochemical reactions. Although NOMs can produces some reactive species on irradiation, its inner filter effect and radical scavenger effects are significantly high and likely to decrease the rate of pollutant degradation in practical applications. To test this hypothesis, lindane degradation was carried out in the presence of 1 mg/L of humic acid. Figure 3.19 shows the effects of humic acid on UV/PMS photochemical degradation of lindane. The results showed that the rate of degradation dramatically decreased in the presence of HA. At reaction time of 3h, the degradation efficiency of lindane decreased from 99 to 70%, when 1.0 mg/L humic acid was added to solution. This is because HA

• •− acts as radical scavenger via competing for OH and SO4 . Beltran and co-workers

(1998) have demonstrated the scavenging effects of humic acids on the degradation of nitrobenzene and 2,6-dinitrotoluene. Gogate and Pandit also concluded that high concentration of humic acid can act as strong scavenger for hydroxyl radicals (2004).

Guan and co-workers (2013) reported that the degradation efficiency of atrazine decreased from 98% to 23% when 3.2 mg/L of NOM was added to the solution. Chu and co-workers (2011) demonstrated that presence of NOMs, especially the humic acid, will

2− quench the radicals in the UV/S2O8 /H2O2 process. Therefore, pre-treatment caution should be considered to reduce the retardation effect due to the presence of non-target compounds, such as HA in real applications.

142

pH = 8.0 pH = 4.0 1.0 pH = 6.0

0.8

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Fig. 3.18. Effect of solution pH on UV/PMS photodegradation of lindane. Experimental conditions: [lindane]̥0 = 3.43 μM and [PMS]̥0 = 300 μM.

1.0 1 mg/L humic acid Blank

0.8

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Figure 3.19. Decomposition of lindane as a function of time in the presence of humic acid. Experimental conditions: [PMS] = 300 μM; [Humic acid] = 1 mg/L and

[lindane]0 = 3.43 μM. 143

3.4.6.4 Effect of inorganic anions on the UV/PMS process

Since real waters usually contain inorganic ions coexisting with organic pollutants, the effect of inorganic anions on the degradation of lindane was studied at the anions

2− − initial concentrations of 1mM. The inorganic ions investigated include CO3 , HCO3 ,

2− − SO4 and Cl , which were added as sodium salts of these anions. Under the ion free condition, 99% lindane was removed at 3 hour reaction time. In the presence of various ions, the percent removal efficiency of lindane was changed to different levels as shown in Figure 3.20. The results showed that the rate of degradation was greatly decreased with

− 2− − 2− the addition of HCO3 and CO3 ions. The presence of Cl and SO4 ions only slightly decreased the degradation efficiency of the process. The effects of Cl− on the degradation mechanism can be different under various oxidation conditions. Several researchers reported that Cl− can have negative effects on pollutant removal efficiency (Ocampo-

Perez et al., 2011). However, Wiszniowski and co-workers (2003) have shown that the addition of Cl− did not have any influence on the mineralization of organic matters. Wang and co-workers (2011) recently reported that the rate of Acid Orange 7 is increased in the

− − presence of Cl ions. Maurino and co-workers (1997) reported, the presence of HCO3

2˗ has strong negative effect on the efficiency of the H2O2/UV and S2O8 /UV processes.

The inhibitory effects of these anions on lindane degradation is mainly attributed to the

•− scavenging of reactive radicals (SO4 ) by these ions, leading to the formation of less

•− reactive radicals. The reactions of these ions with SO4 are given below (Spinks and

Woods, 1990, Huie and Clifton, 1990; Padmaja et al., 1993; Wine et al., 1989).

2‒ •− •− 2‒ 6 −1 −1 CO3 + SO4 → CO3 + SO4 (k = 4.1 × 10 M s ) (3.55)

‒ •− •− 2‒ + 6 −1 −1 HCO3 + SO4 → CO3 + SO4 + H (k = 2.8 × 10 M s ) (3.56) 144

‒ •− • 2‒ 8 −1 −1 Cl + SO4 → Cl + SO4 (k = 2.6 × 10 M s ) (3.57)

2− • •− − SO4 + OH → SO4 + OH (3.58)

Photochemical removal rate of lindane was modified in the following order: blank

2− − 2− − > SO4 > Cl > CO3 > HCO3 . The results showed that photochemical reaction was

2− − 2− rigorously inhibited by CO3 and HCO3 , while it was lightly affected by the SO4 and

Cl− ions. At any rate, the enhanced negative effects of these ions on the photochemical decomposition of lindane would be considered in real water samples as these ions are ranked among the commonest anions.

145

1.0 No additive 2- SO4 Cl- CO 2- 0.8 3 - HCO3

0.6

(lindane)

C/C 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Fig. 3.20. Decomposition of lindane as a function of time in the presence of different inorganic anions. Experimental conditions: [PMS] = 300 μM; [Additive ions] = 1000 μM, [lindane] = 3.43 μM.

146

3.4.7 Mineralization study

It has been reported that some of the intermediate and final by-products generated during degradation process can be more toxic than the parent organic compounds (Idaka et al., 1987, Sweeney et al., 1994). Mineralization of pollutants into harmless inorganic or organic constituents is usually an important parameter in the ultimate removal of toxic compounds in advanced waste water treatment. Mineralization studies were carried out by evaluating the removal of total organic carbon (TOC) in the lindane solutions at different time intervals. Figure 3.21 shows the results of the TOC removal in the lindane solutions (C0= 17.15 μM) in different photochemical processes using PMS as oxidant. As can be seen in Figure 3.21, the highest TOC removal efficiency (92%) was achieved in the UV/Fe2+/PMS system. The main reason is that UV/Fe2+/PMS system provides the

•− • most favorable conditions for the generation of SO4 and OH radicals in the system. The

•− • 2+ overall production and destruction of SO4 and OH radicals by the UV/Fe /PMS system

•− • has already been discussed in the previous sections. The SO4 and OH radicals produced in the UV/Fe2+/PMS system are ultimately involved in the oxidation and mineralization of organic pollutants.

In the present study, varying amount of Fe2+ and PMS were employed and effect of Fe2+ and PMS concentrations on the TOC removal efficiency of the UV/Fe2+/PMS system was studied. It was found out that the efficiency of TOC removal is greatly influenced by the concentration of both Fe2+ and PMS. It is found that the TOC removal efficiency of the system was increased with the increase of PMS concentration. This trend has also been reported by other researchers (Pagano et al., 2011). The improved

TOC removal efficiency at the higher PMS concentration is due to the increasing number

147

−• of SO4 radicals produced under the condition of higher PMS concentration. Three different PMS concentration i.e. 100, 300 and 500 µM were applied and the TOC removal efficiency was determined, as presented in Table 3.11.

To test the effect of Fe2+ concentration, two different concentrations of Fe2+ i.e.

50 and 500 µM were applied for a fixed amount of PMS (500 µM) and the TOC removal efficiency was determined, as presented in Figure 3.22. As can be seen in Figure 3.22, the

TOC removal efficiency was found to decrease with the increase of Fe2+ concentration. It is found that 92% and 88% TOC removal is achieved in 180 min, when the Fe2+ concentration is 50 and 500 µM, respectively. Thus using 500 µM PMS and 50 µM Fe2+,

92% TOC removal is achieved in the UV/Fe2+/PMS system when the pollutant/catalyst/oxidant molar ratio is 1:3:30. The decrease in TOC removal efficiency

2+ •− with increasing Fe concentration can be attributed to enhanced scavenging of SO4 radicals by Fe2+ beyond the optimum concentration limit. It can be concluded from the mineralization results that moderate concentration of PMS (500 μM) can effectively remove the TOC in lindane solution in the UV/Fe2+/PMS system by using micromolar quantities of Fe2+ catalyst.

The complete oxidation of lindane to inorganic constituents, like carbon dioxide

− (CO2) and chloride (Cl ) is represented stoichiometrically as:

•− 2− − + 24SO4 + C6H6Cl6 + 12 H2O → 24 SO4 + 6CO2 + 6Cl + 30H (3.59)

•− Thus, 24 mol of SO4 is required for complete oxidation of 1 mol of lindane. For

•− the conditions of this study, the theoretical need of SO4 with 0.0172 mM lindane is

•− 0.412 mM, however, 0.500 mM SO4 was used in actual experiments. About 19 % loss

•− in the concentration of SO4 occurred in the course of lindane mineralization in the

148

UV/Fe2+/PMS system. The various intermediates formed from lindane decomposition

•− •− may be responsible for loss of SO4 radicals. A part of SO4 radicals may be

•− − decomposed by the undesirable reactions of SO4 with the PMS (HSO5 ) itself and its reaction with the Fe2+ ions.

Comparing the TOC removal efficiencies of the various systems given in Figure

3.21, it can be seen that UV assisted PMS processes (UV/PMS and UV/Fe2+/PMS) always exhibited higher TOC removal efficiency as compared to the Fe2+ activated PMS processes (Fe2+/PMS and Fe2+/PMS/tube-light). Several reasons can be assigned to this result. From foregoing discussion and also from the literature reports, it is known that UV

2+ •− • radiation and Fe catalyst both can activate PMS to give SO4 and OH radicals.

However, the efficiency of the two sources i.e. UV radiation and Fe2+ catalyst in activating the PMS is different from each other (Antoniou et al., 2010a). Anipsitakis and

Dionysiou (2004b) studied the efficiency of three common oxidants (PS, PMS and H2O2) for water decontamination under the influence of transition metals and UV radiation and it was found that UV radiation can more efficiently decompose PMS as compared to the

2+ •− transition metal ions. Also, in the case of Fe /PMS system, significant amount of SO4 is scavenged by the Fe2+ itself, thus lowering the TOC removal efficiency in lindane

2+ •− solution. The results also show that Fe ions can scavenge the SO4 more effectively

•− 2+ when SO4 radicals are generated in slow manner like in the case of Fe /PMS and tube- light/ Fe2+/PMS as compared to the UV/Fe2+/PMS system. From the above discussion it can be concluded that the UV/Fe2+/PMS system is the most efficient method for rapid

•− mineralization of lindane (92% TOC reduction) compared to the other SO4 radical based systems. The increased TOC removal efficiency in the case of UV/Fe2+/PMS

149

system is obviously due to the efficient activation of PMS by the combined effects of UV light and Fe2+ catalyst. Similarly, the regeneration of Fe2+ catalyst (equation 3.11) by UV radiation can also contribute to enhanced mineralization in the UV/Fe2+/PMS system

(Faust and Hoigné , 1990; Zhao et al., 2004). It is found that the TOC removal efficiency of the UV/Fe2+/PMS system decreased with the passage of time. The observed decrease in the TOC removal efficiency with the reaction time can be attributed to the formation and accumulation of stable organic acids, which appear in the degradation progresses of several advanced oxidation processes. The mineralization studies revealed that although

11 and 40% lindane is degraded in the Fe2+/PMS and tube-light/ Fe2+/PMS systems, respectively in 180 min, the TOC removal efficiency was less than 10% in each case.

This means that there is time span between degradation and mineralization of lindane.

It can be summarized that 6, 9, 36 and 92% TOC removal is achieved in the

Fe2+/PMS, tube-light/Fe2+/PMS, UV/PMS and UV/Fe2+/PMS systems, respectively in 3 hours, using 500 μM of PMS. Thus the order of TOC removal efficiency for different systems is: UV/Fe2+/PMS > UV/PMS > tube-light/Fe2+/PMS > Fe2+/PMS.

150

1.0

0.8

0 0.6

Fe2+/PMS

TOC/TOC 0.4 Tube-light/Fe2+/PMS UV/PMS UV/Fe2+/PMS 0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Reaction time (min)

Figure 3.21. Removal of TOC as a function of time in different processes: (a) Fe2+/PMS, (b) Tube-light/Fe2+/PMS, (c) UV/PMS and (d) UV/Fe2+/PMS systems.

Experimental conditions: [lindane]0 = 17.15 μM, [PMS]0 = 500 μM in all cases, 2+ 2+ [Fe ]0 = 500 μM for process (a) and (b) and [Fe ]0 = 50 μM for process (d), pH = 3.0 for process (a), (b) and (d), and pH = 6.0 for process (c).

2+ 1.0 [Fe ] = 500 M [Fe2+] = 50 M

0.8

0 0.6

TOC/TOC 0.4

0.2

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Figure 3.22. Effect of Fe2+ concentration on TOC removal in the UV/Fe2+/PMS system.

Experimental conditions: [lindane]0 = 17.15 μM, [PMS]0 = 500 μM, pH = 3.0.

151

Table 3.11. TOC removal efficiency measured under different initial concentration of

PMS.

Time (min) TOC removal (%)

([PMS]0 = 100 µM) ([PMS]0 = 300 µM) ([PMS]0 = 500 µM)

0 0 0 0

40 4.3 16 19

80 6.8 35 42

120 15 56 88

180 22.7 64 92

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3.4.8 Oxidant residue analysis of UV/Fe2+/PMS process

The remaining amount of PMS with the corresponding decrease in TOC removal in the UV/Fe2+/PMS system is shown in Figure 3.23. The reaction was started with different initial concentrations of PMS (100, 300 and 500 μM) and the TOC removal efficiencies were determined. It can be seen that the amount of PMS oxidant decreases as the TOC removal proceeds with the reaction time. More than the stoichiometric amount of PMS or any other oxidant is generally used in the mineralization of pollutant.

Considerable amount of PMS is generally consumed in the degradation of the intermediate by-products. From Figure 3.23, it can be inferred that the oxidant decomposition profiles differ from the TOC removal profiles. The plausible reason is that

•− • the SO4 and OH radicals may undergo strong radical-radical recombination besides

•− • their reactions with lindane. The undesired recombination reactions of the SO4 and OH species can be significantly minimized when different components of the system such as,

PMS, Fe2+ and lindane are mixed in appropriate ratio.

The stoichiometry of lindane mineralization by PMS oxidant is given in equation

(3.62). The equation (3.62) shows that a total of 0.412 mM of PMS is required for complete mineralization of 0.0172 mM of lindane. The results of the PMS residue analysis revealed that 500 µM of PMS is consumed in 180 min and the TOC removal efficiency is only 92%. The apparent difference (19%) between the stoichiometric and the experimental amount of PMS can be attributed to the existence of some stable intermediates generated during the course of lindane mineralization. Thus excess PMS is needed for decomposition of these intermediate by-products in addition to the amount required for the target pollutant. Also, some amount of PMS is lost due to the undesired

153

reaction of the PMS with itself and with the Fe2+ ions, as shown in equations (3.3) and

(3.4).

One of the main aims of the PMS residue analysis was to evaluate if the amount of the PMS applied to the system is totally consumed during the reaction or it may remain in the system after the degradation of the pollutant. From the simultaneous study of the

PMS residues and the TOC removal, it is easy to device optimum doses of PMS for decomposing specific amount of pollutant. The PMS residue analysis also provides useful information about the rate of decomposition and consumption of PMS oxidant. From the

Figure 3.23 it can be seen that PMS is readily decomposed and consumed in the

UV/Fe2+/PMS system.

154

Time of photolysis (min)

400 300 200 100 0

1.0 [PMS] = 100 M 1.0 [PMS] = 300 M [PMS] = 500 M 0.8 0.8

0 0.6 0.6 0

0.4 0.4

TOC/TOC

PMS/PMS

0.2 0.2

0.0 0.0

0 100 200 300 400

Time of photolysis (min)

Figure 3.23. Residues of PMS with subsequent TOC removal as a function of time in the UV/Fe2+/PMS process at varied PMS concentrations. Experimental conditions:

2+ [lindane]̥o = 17.15 μM, [Fe ]̥o = 50 μM and pH = 3.0.

155

3.4.9 By-product analysis and reaction mechanism of the UV/PMS system

To There were mainly six reaction by-products identified by GC/MS in this study, i.e., 1,1,2,3,4,5,6-heptachlorocyclohexane (HeCH), 1,2,3,4,5,6-hexachlorobenzene

(HCB), 1,3,4,5,6-pentachlorocyclohexene (PCCH), 3,4,5,6-tetrachlorocyclohexene

(TeCCH), 1,2,4-trichlorobenzene (TCB) and 2,4,5-trichlorophenol (TCP). The identified by-products, along their molecular formula, molecular weight and chemical structure, are shown in Table 3.12. These by-products have been reported previously in various oxidative studies on lindane, e.g., HeCH, HCB, PCCH, TeCCH and TCP in POMs photocatalysis (Antonaraki et al., 2010), TCB and HCB in the photo-Fenton reaction

(Nitoi et al., 2013) as well as HeCH, PCCH and TeCCH in TiO2 photocatalysis (Zaleska et al., 1999). In this study, the exact reacting radical could not be distinguished with

• •− certainty because of the coexistence of both OH and SO4 in the reaction solution (Khan et al., 2014, Antoniou et al., 2010a). Due to the similarities in the reaction mechanism of these two radicals, it is very likely that the detected by-products can be from either radical reaction. A potential reaction pathway for the degradation of lindane was proposed and is shown in Figure 3.24, including (1) dechlorination, (2) chlorination, (3) dehydrogenation and (4) hydroxylation.

(1) Dechlorination reaction is presumably resulted from homolytic scission of the

C-Cl bond upon UV excitation (Guillard et al., 1996, Legrini et al., 1993, Somasundaram and Coats, 1991). As a result, chlorine (Cl•) is released from lindane, leaving a carbon centered radical. A subsequent abstraction of hydrogen from adjoining carbon atom by

• •− OH and/or SO4 might lead to the formation of PCCH. A sequential loss of further

156

chlorine resulted in the formation of lesser chlorinated by-products such as TeCCH and

TCB.

• •− • (2) Chlorination of lindane can be resulted from its reaction with Cl or Cl2 . Cl

•− 0 and Cl2 are strong oxidizing species with E of 2.4 and 2.0 V, respectively (Alegre et al., 2000). These oxidizing species may react with organic compounds via addition to double bond, hydrogen abstraction or electron-transfer reactions (Gilbert et al., 1988).

Thus an abstraction of hydrogen from lindane via Cl• may lead to the formation of HeCH

(reaction (3.60)). A similar explanation has been provided by Antonaraki et al.

(Antonaraki et al. 2010) employing POMs photocatalysis. In fact, the interaction of Cl−

•− with SO4 with the subsequent formation of the chlorinated organic compounds in water has been commonly reported in literature (Anipsitakis et al., 2006).

(3.60)

(3) Dehydrogenation may occur with the abstraction of two adjoining hydrogen

• •− atoms by OH and/or SO4 attack. The formation of a stable HCB by-product allowed the formation of a double bond to be thermal kinetically possible (Guo et al. 2000). The formation of HCB from lindane via hydrogen abstraction by •OH is also reported elsewhere (Antonaraki et al. 2010).

(4) Hydroxylation is a process that introduces a hydroxyl group (─OH) into an organic compound. The addition of the electrophilic •OH to TCB forms a carbon centered

• radical, which by addition of O2 yields a peroxy radical. After releasing HO2 , TCP could

157

be formed (reaction (3.61)). Hydroxylation of chlorobenzene and other chlorinated aromatic compound through such a pathway has been proposed earlier (Drijvers et al.

1998, Zona et al. 2002).

(3.61)

•− •− In SO4 mediated mechanisms, SO4 oxidizes the aromatic ring to a radical cation, which upon hydrolysis leads to the formation of hydroxycyclohexadienyl radical.

• The resulting radical, after reaction with O2 and releasing subsequently HO2 , is converted into a hydroxylated phenolic by-product (reaction (3.62)) (Neta et al. 1977,

Anipsitakis et al., 2006, Walling and Camaioni, 1975).

Though not identified by our method, ring opening and cleavage by-products are also expected to be formed. The intermediate by-products HCB, TCB and TCP, for

− example, were known to mineralize into CO2, H2O and Cl with an extended reaction time by photocatalytic and photochemical transformations (Hiskia et al. 2000, Lin et al.

2011, Wang and Chu, 2013).

158

Figure 3.24. Proposed degradation pathway of lindane by UV/PMS. [Lindane]0 = 17.15

μM, [PMS]0 = 500 μM, pH = 5.8. 159

Table 3.12. Reaction intermediates of lindane along with chemical structure, molecular formula and molecular weight.

Compound name Molecular formula Molecular weight Chemical structure

1,1,2,3,4,5,6-

Heptachlorocyclohexane C H Cl 325 6 5 7 (HeCH)

Lindane C6H6Cl6 291 (γ-HCH)

1,2,3,4,5,6-

Hexachlorobenzene C Cl 285 6 6 (HCB)

1,3,4,5,6-

Pentachlorocyclohexene C H Cl 254 6 5 5 (PCCH)

3,4,5,6- tetrachlorocyclohexene C H Cl 219 6 6 4 (TeCCH)

2,4,5-Trichlorophenol

(TCP) C H Cl OH 197 6 2 3

160

(Table 3.12. continue)

Compound name Molecular Molecular weight Chemical structure

formula

1,2,4-Trichlorobenzene

(TCB) C H Cl 181 6 3 3

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3.4.10 Degradation of Trichlorobenzene

Trichlorobenzene (TCB) is one of the most commonly reported intermediate generated during the degradation of lindane by different techniques (Cristol 1947; Orloff

1954; Li et al., 2011). In the current study, TCB was found as a major intermediate byproducts identified in the UV/PMS and UV/Fe2+/PMS processes. Some of the important isomers of trichlorobenzene (TCB) are: 1,2,4-TCB, 1,3,4 and1,3,5-TCB. The three isomers of TCB can be produced in varying amount in different degradation processes. In our study, all the three isomers were produced; however, the quantitative amount of the different isomers was not determined, separately.

TCB are colourless liquids with a pleasant smell. They are only slightly soluble in water, but easily soluble in organic solvents. They are non-flammable and they can easily decompose to produce toxic gases when heated. TCBs belong to the group of compounds, commonly known as the volatile organic compounds (VOCs). The main use of TCBs is in the chemical industry: in the manufacture of dyestuffs and textiles, and in synthetic oils. They are also used as lubricants and heat transfer fluids, as wood preservatives, as cleaning agents for septic tanks and in abrasive formulations. TCBs were used in tropical regions of the world as an insecticide against termites. In the past,

TCBs were also formed during the manufacturing of lindane as a by-product. TCBs can also be produced as a by-product during industrial cracking (Hooftman and de Kreuk,

1982) and other combustion processes (Jay and Stieglitz, 1995; Panagiotou et al., 1996).

In this part of the study, the efficiency of various UV/oxidant processes, such as

UV/H2O2, UV/PMS and UV/PS for the removal of aqueous TCB were investigated. With batch experiments, 3.43 μM aqueous TCB solutions were treated in different

162

UV/peroxides systems and the TCB removal efficiency were determined. The TCB readily undergoes decomposition in the UV/oxidant systems. It is found that all the above mentioned UV/oxidants systems can effectively degrade TCB. Using the different

UV/oxidant systems, more than 90% TCB was degraded in less than half hour reaction time. Thus all the three UV/oxidants systems gave good results and the degradation efficiency varied in the order; UV/H2O2 > UV/PMS > UV/PS. On UV photolysis, H2O2

• •− • generate two OH radicals, PMS generates one SO4 and one OH radical, while PS

•− generate two SO4 radicals. Thus H2O2 and PMS based processes involve the generation

• •− • of OH radical, while PS process involve only SO4 radicals generation. OH radical has strong affinity for conjugated systems and this seems to be the main reason which is responsible for higher TCB removal efficiency by the UV/H2O2 system as compared to the UV/PMS and UV/PS systems.

Pseudo-first order kinetics was used to describe the TCB degradation by various

UV/oxidants systems. Under first order kinetics the rate of decomposition can be expressed by equation 3.54: (i.e., −ln[C]/[C0] = kt) where Co: initial concentration of

TCB (µM/L); C: concentration of TCB (µM/L) at time t; t: reaction time (min); k: pseudo-first order rate constant (min–1).

Based on the plots of −ln[C/C0] versus time (t), a linear relationship was established and it was used to calculate the degradation rate constant ‘k’ from a linear least-squares fit of the experimental data. In the current study, different initial concentrations of the various oxidants were applied and efficiency of the photochemical degradation of TCB was evaluated based on the pseudo-first order reaction rate constants

‘k’. The result of TCB removal by various UV/oxidant is shown in Figure 3.25a, 3.25b

163

and 3.25c. The results showed that for a given concentration of TCB, the rate of TCB degradation increased with increasing the initial concentration of the oxidant. The enhanced degradation efficiency of the UV/oxidants processes with the increase in oxidant concentration is due to increasing number of reactive radicals (•OH and/or SO•−).

Under the optimum conditions, using 100 μM of oxidant, 98, 95 and 70% TCB was removed in the UV/H2O2, UV/PMS and UV/PS processes, respectively in 25 min. The overall production and consumption of •OH and SO•− radicals that can take place in the

UV/H2O2, UV/PMS and UV/PS systems is discussed in detail in the previous sections.

It can be concluded from the results that the reaction follows pseudo-first order kinetics in all cases. However, the rate of reaction was found to vary with the nature of the oxidant. In Figure 3.26, comparison of TCB degradation by various UV/oxidants systems including UV/H2O2, UV/PMS and UV/PS is shown. The overall order of the various processes is: UV/H2O2 > UV/PMS > UV/PS.

Some useful information can be collected from the degradation of TCB in the

UV/peroxides processes. As compared to the parent compound i.e. lindane, it is seen that trichlorobenzene (TCB) undergoes faster degradation in the UV/H2O2 process rather than

UV/PMS or UV/PS process. Also, the overall degradation of TCB by UV/peroxides processes is faster than that of lindane. This is probably due to the presence of unsaturated bond inside the TCB molecule, that is chemically more reactive towards oxidation by •OH radical as compared to non-conjugated compounds, such as lindane.

Under the conditions of UV/oxidants alone, TCB undergoes 100% degradation in reaction time of 40 minutes. On the other hand, less than 40% lindane degradation was observed when similar experimental conditions were provided. The observed degradation

164

rate constants values ‘k’ calculated from the plots of −ln(C/C0) vs. reaction time were

−1 −1 −1 −1 found to be 3.04 × 10 , 2.85 × 10 and 1.50 × 10 min in the UV/H2O2, UV/PMS and

UV/PS processes, respectively.

165

(a)

[H2O2] = 0.1 (mM) 1.0 5 [H2O2] = 0.3 (mM)

[H2O2] = 0.5 (mM) 4

0.8 3

)

0 2

0.6 -ln(C/C 1

(TCB)

0 0

C/C 0.4 0 5 10 15 20 25 30 Time of photolysis (min)

0.2

0.0 0 10 20 30 40 50 Time of photolysis (min)

(b)

7 [PMS] = 0.1 mM 1.0 [PMS] = 0.3 mM 6 [PMS] = 0.5 mM 5

0.8 ) 4

0 3

0.6 -ln(C/C 2

(TCB) 0 1

C/C 0.4 0 0 10 20 30 40 50 Time of photolysis (min) 0.2

0.0 0 10 20 30 40 50 Time of photolysis (min)

166

(c)

5 [PS]=0.1 mM 1.0 [PS]=0.3 mM 4 [PS]=0.5 mM 3 0.8

) 0 2

1 0.6 -ln(C/C

(TCB) 0 0

C/C 0.4 0 10 20 30 40 50 Time of photolysis (min)

0.2

0.0 0 10 20 30 40 50 Time of photolysis (min)

Figure 3.25. Trichlorobenzene (TCB) decay as a function of UV photolysis time in the;

(a) UV/ H2O2, (b) UV/PMS, (c) UV/PS system. Experimental conditions: [TCB]̥0

= 3.43 μM, [H2O2]̥0 = [PMS]̥0 = [PS]̥0 = 0.1-0.5 mM, pH = 5.8. In the inset, −ln(C/C0) Vs time of photolysis is plotted for determination of rate constant, k.

5 UV/PS 1.0 UV/PMS 4 UV/H2O2

0.8 3

) 0 2 0.6

-ln(C/C

(TCB) 1

C/C 0.4 0

0 5 10 15 20 25 30 Time of photolysis (min) 0.2

0.0 0 10 20 30 40 50 Time of photolysis (min)

Fig. 3.26. Comparison of the decay of TCB by PS/UV, H2O2/UV and PMS/UV

processes. Experimental conditions: [TCB]0 = 3.43 μM, Initial oxidant concentration = 300 μM, pH = 5.8. 167

3.5 Hydroxyl radical based AOP using UV/H2O2 system for lindane degradation

H2O2 (HOOH) is conventionally well known inorganic peroxide and strong oxidizer with standard oxidation–reduction potential (E0) of 1.776V (Betterton and

Hoffmann 1990). However, the chemical is very stable which can react with substrates at moderately slow rate. Several kinds of activation sources such as UV radiation, gamma rays and/or transition metal ions are commonly used as activator of H2O2 oxidant. The most common AOPs based on H2O2 include Fenton’s reagent (Burbano et al., 2008;

Elmolla and Chaudhuri, 2009) and UV/H2O2 (Xu et al., 2009; Kalsoom et al., 2012).

Several useful reviews have been published on the application of H2O2 in wastewater treatment (Neyens and Baeyens, 2003; Gogate and Pandit, 2004).

Highly reactive and non-selective hydroxyl radicals (•OH, E˚ = +1.8 to +2.7V) can be generated by the activation of H2O2 with UV radiation via equation 3.60 (φ=1.0).

Those processes in which the concentration of target contaminant changes with time, while the concentration of oxidant remains relatively constant, are generally described by pseudo-first order kinetics. Under first order kinetics, the rate of degradation is directly proportional to the sum of the material, undecomposed as mathematically given in equation 3.54. Equation 3.54 is generally used to calculate the degradation rate constants, k.

Degradations rate constants ‘k’ of lindane and t1/2 described by use of pseudo-first order kinetics at using different initial concentration of lindane are presented in Table

3.13. The results revealed that the pseudo-first order reaction rate constants decreased with the increasing initial concentration of lindane. This is in accordance with the results observed by Danishvar and co-workers, which indicated that the rate of nitro-phenol

168

degradation increased with the decreasing pollutant concentration in the UV/oxidant process (Daneshvar et al., 2007).

In the current study, different initial concentrations of H2O2 were applied and efficiency of the photochemical degradation of lindane was evaluated based on the pseudo-first order reaction rate constants ‘k’. The result of lindane removal by the

UV/H2O2 system is shown in Figure 3.27a and the effect of initial H2O2 concentration on the degradation rate constant ‘k’ is shown in Figure 3.27b. Table 3.14 gives the pseudo- first order rate constants, k for degradation of lindane at different initial concentration of

H2O2. In Figure 3.28, the variation of observed degradation rate constant, k with increasing concentration of H2O2 is shown. The results showed that for a given concentration of lindane, the rate of lindane degradation increased with increasing the initial concentration of H2O2. The results further showed that the increase in degradation rate with the increasing H2O2 concentration is fast in the beginning, while it slows down with time until the rise becomes very slight after some optimum concentration of H2O2.

The enhanced degradation efficiency of the UV/H2O2 process with the increase in H2O2 concentration is due to increasing number of •OH produced at higher concentration of

H2O2. However, this factor becomes less obvious at the stages when the recombination

• • reactions among the OH and the reaction of OH with H2O2 molecules surmount the reaction of •OH with lindane. Such an increase in the degradation rate with the increasing

H2O2 concentration is observed in the UV/H2O2 oxidation of several other pollutants

(Behnajady et al., 2004; Riga et al., 2007). Under the optimum conditions, using 500 μM

H2O2, 49% lindane was removed in 180 min, when the lindane/H2O2 molar ratio is 1:87.

169

• The overall production and consumption of OH taking place in the UV/H2O2 process can be given in the following reactions (Christensen et al., 1982, Buxton et al.,

1988, Chen et al., 2011, Khan et al., 2013);

• −1 H2O2 + hv → 2 OH (Φ= 1.0 mol einstein ) (3.61)

• • 9 −1 −1 OH + H2O2 → HO2 + H2O, (k = 2.7 x 10 M s ) (3.62)

• 2 OH → H2O2 (3.63)

− + H2O2 → HO2 + H (3.64)

• − •− 9 −1 − OH + HO2 → H2O + O2 , k= 7.5 x 10 M s (3.65)

•− − • O2 + H2O2 → OH + OH (3.66)

• • 10 −1 −1 OH + HO2 → H2O + O2, k = 1 x 10 M s (3.67)

− + H2O → HO + H (3.68)

3.5.1 Comparison of the UV/PMS system with UV/H2O2 process

The two peroxides mentioned above (PMS and H2O2) are similar in structure and

− both contain O–O bond. One hydrogen atom in H2O2 (HOOH) is replaced by SO3 to

− − generate HOOSO3 (PMS). Due to the influence of the SO3 , O–O bond is lengthened

− − and the bond energy decreases. The distances of the O–O bonds in HSO5 (HOOSO3 ) and solid H2O2 are 1.453 and 1.460 Å, respectively (Flanagan et al., 1984). The estimated bond energy in H2O2 is 213.3 kJ/mol, while it is estimated that PMS is less in bond energy than H2O2 (Reints et al., 2000, Yang et al., 2010). Due to increased O-O bond length in PMS, it is activated more effectively by UV radiation than the H2O2 oxidant.

Control experiments were conducted to determine the lindane removal by solely H2O2.

The results showed that the H2O2 oxidant alone did not give any significant degradation of lindane within 3 hour reaction (data not shown).

170

It can be concluded from the results that the reaction follows pseudo-first order kinetics in each case. The rate of reaction was found to vary considerably with the change of the oxidant types such that: UV/PMS > UV/H2O2. Our results are in accordance with literature reports for similar reactions. Antoniou and co-workers (2010a) reported that

UV/PMS system has greater efficiency than UV/H2O2 system in the removal of microcystin-LR.

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Table 3.13. Kinetic data of lindane photodegradation with UV/H2O2 ([H2O2]0 = 300 µM) processes at different initial concentration of lindane. ______−1 2 C0 (µM) kobs (min ) t1/2 (min) R value

0.68 4.07 × 10˗3 170 0.996

3.4 2.85 × 10˗3 243 0.984

6.8 1.22 × 10˗3 568 0.957

Table 3.14. Pseudo-first order rate constants for degradation of lindane (3.43 µM) by UV photolysis in the presence of various amounts of hydrogen peroxide. −1 2 No. [H2O2]0 (µM) kap (min ) R

1 100 2.08 × 10−3 0.955

2 300 2.85 × 10−3 0.984

3 500 3.71 × 10−3 0.946

4 1000 4.14 × 10−3 0.965

5 2000 4.56 × 10−3 0.966

172

(a)

1.0

0.8

0.6

(lindane)

C/C 0.4 [H2O2] = 0.1 mM

[H2O2] = 0.3 mM [H O ] = 0.5 mM 0.2 2 2 [H2O2] = 1.0 mM

[H2O2] = 2.0 mM

0.0 0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

(b)

1.0 2 [H2O2] = 2.0 mM, R = 0.9868 2 [H2O2] = 1.0 mM, R = 0.9829 0.8 2 [H2O2] = 0.5 mM, R = 0.9573 2 [H2O2] = 0.3 mM, R = 0.9844 2 [H2O2] = 0.1 mM, R = 0.9672 0.6

)

-ln(C/C 0.4

0.2

0.0

0 20 40 60 80 100 120 140 160 180 200 Time of photolysis (min)

Figure 3.27. (a) Photochemical decay of lindane under different initial concentrations of

H2O2, (b) plots of –ln(C/C0) Vs time of photolysis for different initial

concentrations of H2O2. Experimental conditions: [lindane]̥o = 3.43 μM, pH = 6.0. 173

0.0050

0.0045

)

-1 0.0040

0.0035

0.0030

0.0025

Rate constant, k (min constant, Rate

0.0020

0.0015 0.0 0.5 1.0 1.5 2.0 2.5 Concentration of H O (mM) 2 2

Figure 3.28. Variation of observed degradation rate constant, k with change in concentration of H2O2. Experimental conditions: [lindane]̥0 = 3.43 μM, and pH = 6.0.

174

3.6 Photocatalytic activity of sulfur doped TiO2 (S-TiO2) and TiO2 under Visible light

Figure 3.29 shows the visible light activity (VLA) of S-TiO2 for the degradation of lindane in aqueous solution. The results of control experiments, including (i) visible light only, (ii) ref-TiO2/dark, (iii) S-TiO2/dark, and (iv) ref-TiO2/visible light, revealed that neither activation of ref-TiO2 (band gap energy, EG = 3.18 eV) (Han et al., 2011), nor direct photodegradation of lindane with visible light (λ > 420 nm) was effective in this study. The results; however, showed that significant degradation of lindane occurred in visible light-assisted S-TiO2 photocatalysis (S-TiO2/vis), leading to 31.0% lindane removal in 6 hr. The VLA of S-TiO2 is associated with its reduced band gap value (i.e.,

2− 2.94 eV), induced by substitutional doping of S in the TiO2 lattice (Han et al., 2011).

Consequently, the absorption edge of S-TiO2 was shifted to lower energy region, thereby capable of absorbing visible light photon for the promotion of electrons to the conduction band (Han et al., 2011, Umebayashi et al., 2002). The photogenerated electron (e−)

(reaction 4.75) has high reduction potential, capable of reducing surface adsorbed oxygen

•− (O2), and yielding superoxide radical anion (O2 ) (reaction (4.79)) (Zhao et al., 2014,

+ Banerjee et al., 2014). Contrary to UV/TiO2 process, the photogenerated hole (h )

− resulted in visible light-assisted S-TiO2 photocatalysis cannot oxidize H2O or HO , because of thermodynamics unsuitability, thus avoiding the formation of •OH (Zhao et al., 2014, Banerjee et al., 2014). However, the formation of •OH in visible light-assisted

•− S-TiO2 photocatalysis is reported to have taken place via O2 pathways (reactions (4.80)-

(4.84)) (Goldstein et al., 2008, Zhao et al., 2014). Thus the photogenerated electrons (e−)

•− • and the associated ROSs like O2 and OH are most likely responsible for the visible light-assisted S-TiO2 photocatalytic degradation of lindane. Though mechanisms of

175

visible light assisted doped-TiO2 photocatalysis are not currently well-established, we will attempt later, to identify the role of various reactive oxygen species (ROSs) generated in S-TiO2/vis process, based on the reaction intermediates identified via

GC/MS analysis.

+ − TiO2 + hν → hVB + eCB (4.75)

+ − hVB + eCB → Energy (4.76)

+ • hVB + H2O → OH (4.77)

+ − • hVB + HO → OH (4.78)

− •− eCB + O2 → O2 (4.79)

− •− + eCB + O2 + 2H → H2O2 (4.80)

•− •− + O2 + O2 + 2H → H2O2 + O2 (4.81)

−• + • O2 + H → HO2 (4.82)

• + 2HO2 + 2H → 2H2O2 + O2 (4.83)

− + • eCB + H2O2 + H → H2O + OH (4.84)

Degradation of lindane by the visible light-induced S-TiO2 photocatalysis followed pseudo first-order kinetics, as shown by following equation (4.85):

C -ln  kobs t (4.85) C0 where C0 and C are the concentration of lindane at initial and after time t, respectively; t is the radiation time; and kobs is the observed pseudo first-order rate constant. The visible light-induced S-TiO2 photocatalysis is effective for degradation of lindane. A further study was performed to investigate the effect of the critical operation parameters such as

176

solution pH, initial concentration of lindane, and catalyst loading on the visible light- induced S-TiO2 photocatalysis for lindane degradation.

177

1.0

0.9

0.8

(lindane)

0 Vis light alone

C/C Ref-TiO2/dark S-TiO2/dark 0.7 Ref-TiO2/vis light S-TiO2/vis light

0.6

0 1 2 3 4 5 6 Irradiation time (h)

Figure 3.29.Photocatalytic degradation of lindane using S-TiO2 photocatalyst under

visible light. Experimental conditions: [lindane]0 = 1.0 µM, [S-TiO2]0 = 0.23 g/L,

pH = 5.8.

178

3.6.1 Factors affecting the efficiency of photocatalytic activity of the S-TiO2

3.6.1.1 Effect of solution pH

The solution pH may affect the surface charge on the photocatalyst and also the state of ionization of the substrate. Thus the adsorption of the substrate as well as its photocatalytic decay is expected to vary with solution pH (Bhatkhande et al., 2002,

Apkan and Hameed, 2009, Evgenidou et al., 2005a). To investigate the influence of solution pH on photocatalytic activity of S-TiO2, three different pH values (i.e., 4.0, 5.8 and 8.0) were chosen for lindane degradation under visible light illumination. The results are shown in Figure 3.30. The results showed that photocatalytic activity of S-TiO2 films varies significantly with the solution pH. The highest lindane removal was observed at pH 5.8, corresponding to 31.0% removal in 6 hr. In contrast, the removal efficiency decreased in stronger acidic as well as basic conditions, indicating 27.2 and 22.4% removal at pH 4.0 and 8.0, respectively, after 6 hr of visible light irradiation. The lindane degradation followed pseudo first-order kinetics model in the studied pH range with kobs of 4.19 × 10−2, 5.93 × 10−2, and 4.14 × 10−2 h−1 for pH 4.0, 5.8, and 8.0, respectively. The obtained results are in good agreement with the literature reports, dealing with photocatalytic degradation of other insecticides such as atrazine and dimethoate (Sacco et al., 2015, Evgenidou et al., 2005a). Since the electrostatic force of attraction or repulsion between the photocatalyst’ surface and the pollutant is affected by the solution pH

(Evgenidou et al., 2005a, Evgenidou et al., 2005b), the degradation efficiency of S-TiO2 photocatalysis was found to change with the varying pH. At the lower pH value, the

−• + concentration of O2 is reduced by the reaction with H ions (reaction (4.82)) (Hoffmann

179

−• at al., 1995), and so the removal efficiency of lindane due to O2 reaction might be decreased.

The removal efficiency of lindane was also reduced at pH 8.0. It can be attributed to the surface charge change of photocatalysts, leading to surface electrostatic repulsion phenomena. Literature studies show that TiO2 photocatalysis of atrazine also resulted in a lower removal efficiency at the increasing solution pH, i.e., 7 or 10 (Parra et al., 2004).

Above all, our results clearly indicate that the highest photocatalytic activity was observed at pH 5.8.

3.6.1.2 Effect of initial concentration of lindane

Figure 3.31 depicts the influence of initial concentrations of lindane on the observed pseudo first-order rate constant, kobs during the S-TiO2 photocatalysis under visible light irradiation. The kobs were found to decrease with increasing the initial concentrations of lindane. Under the experimental conditions in this study, the value of

−2 −2 −2 −1 kobs was 6.50 × 10 , 5.93 × 10 and 5.22 × 10 h for 0.5, 1.0 and 2.0 µM of initial concentration of lindane, respectively. The most plausible reason for decreasing kobs could be the formation of higher concentration of reaction by-products at higher initial concentration of lindane. The produced reaction by-products can compete with lindane molecules for ROS, thus lowering the removal efficiency of lindane (Ghodbane and

Hamdaoui, 2010). Our results are in accordance with the findings of Senthilnathan and

Philip (Senthilnathan and Philip, 2010). The removal efficiency of lindane decreased with increasing initial concentrations of lindane in the solute concentration range of 0.086-

0.517 µM. Wang et al. (Wang et al., 2011) studied the photocatalytic activity of C-N-S-

180

tridoped TiO2 photocatalyst for the removal of tetracycline and found that the removal efficiency decreases at an increasing initial concentration of tetracycline.

However, the initial degradation rate of lindane (calculated by the change in concentration with time at an initial reaction time of 1 hr) was determined at three different initial concentrations (i.e., 0.5, 1.0 and 2.0 µM) and the results are shown in

Figure 3.31. As seen in Figure 3.31, the initial degradation rate of lindane increased since the number of lindane molecules colliding with ROS increased at higher concentrations of lindane (Khan et al., 2013). Under these experimental conditions, the degradation rate of lindane was 5.25 × 10−2, 9.00 × 10−2 and 1.40 × 10−1 µM h−1 at 0.5, 1.0 and 2.0 µM of initial concentration of lindane, respectively.

Figure 3.32 shows the effect of PMS on TiO2 photocatalysis of lindane under visible light irradiation. The results of control experiments showed that only 4.1% lindane was removed by PMS direct oxidation under visible light in 6 hr, indicating that PMS cannot be effectively activated by visible light irradiation (Figure 3.32). The ref-

TiO2/PMS/vis process showed a 7.0% lindane removal in 6 hr, demonstrating that addition of PMS had no significant effect on the efficiency of visible light-assisted ref-

TiO2 photocatalysis of lindane. Interestingly, photocatalytic activity of S-TiO2/vis was significantly increased in the presence of PMS, leading to 68.2% removal of lindane in 6 hr. The enhanced removal efficiency is most probably resulted from the role of PMS as an acceptor of electrons, thereby reducing the rate of the electron-hole recombination

(i.e., opposing reaction (4.76)) on the photocatalyst surface (Hoffmann at al., 1995). Also,

•− • PMS is an efficient source of SO4 and/or OH radicals in the TiO2-based photocatalytic processes, according to the reactions (4.86) and (4.87) (Malato at al., 1998).

181

− − • 2− HSO5 + eCB → OH + SO4 (4.86)

− − − •− HSO5 + eCB → OH + SO4 (4.87)

•− 0 SO4 is a strong oxidant with E of 2.5-3.1 V (Neta at al., 1988), and it exhibits high efficiency on the decomposition of many organic compounds (Anipsitakis and

•− Dionysiou, 2003). Although more selective, SO4 reacts with organic compounds also through some common mechanisms like •OH, i.e., hydrogen abstraction, electron

•− transfer, or addition to double bonds (Khan et al., 2014). Furthermore, SO4 can also trap the photogenerated electrons according to reaction (4.88), thereby reducing the rate of the

•− • electron-hole recombination. Besides, SO4 can interact with water and generate OH, according to reaction (4.89):

•− − 2− SO4 + eCB → SO4 (4.88)

•− • 2− + SO4 + H2O → OH + SO4 + H (4.89)

The effect of initial concentrations of PMS on the photocatalytic activity of S-

TiO2/vis process was investigated, and the resulting kobs are shown in the inset of Figure

3.32. As can be seen, the value of kobs for S-TiO2/vis process increased as the initial concentration of PMS is increased. A plausible expalanation for the increasing kobs could

•− • be the increased concentrations of SO4 and OH via reactions (4.86) and (4.87)), obtained at the increasing concentration of PMS (Khan et al., 2016). As stated earlier, the electron-hole recombination is also inhibited by PMS. So, an increase in the concentration of PMS might cause enhance inhibition of electron-hole recombination, thereby increasing the removal efficiency of lindane, and hence the kobs is increased. The

−1 −1 −1 −1 −1 calculated kobs was found to be 1.35 × 10 , 1.84 × 10 , 3.25 × 10 , and 4.82 × 10 h , when the initial concentration of PMS was 0.1, 0.2, 0.5 and 1.0 mM, respectively. An

182

initial PMS concentration of 0.2 mM was used in all subsequent experiments, which yielded 68.2% lindane removal in 6 hr.

A comparison of TiO2 photocatalysis in terms of percent removal efficiency (%) of lindane was also performed (Table 3.15). The results showed that the percent removal efficiency of lindane followed the order: ref-TiO2/vis < ref-TiO2/PMS /vis < S-TiO2/vis <

S-TiO2/PMS /vis, which corresponded to 4.2, 7.3, 31.0 and 68.2% lindane removal, respectively, in 6 hr. The observed pseudo first-order rate constant (kobs) and the calculated half-life (t1/2) for the processes studied herein are also given in Table 3.15. The results showed that presence of PMS had no apparent effect on the efficiency of ref-

TiO2/vis, indicating that neither visible light nor ref-TiO2 is effective in activating PMS, for lindane degradation. The results given in Table3.15 revealed that the half-life of the various TiO2 photocatalytic processes (except ref-TiO2/PMS/vis) were several fold reduced in the presence of 0.2 mM PMS, thus making visible light-assisted S-TiO2 photocatalysis a suitable option for application purposes. This study showed that addition of PMS is very beneficial to TiO2-based photocatalysis by way of reducing the size of photocatalytic reactor. Consequently, considerably higher removal efficiency of lindane can be achieved in a comparatively less reaction time.

183

1.0

pH0 = 8.0

pH0 = 4.0 pH = 5.8 0.9 0

0.8

(lindane)

C/C

0.7

0.6

0 1 2 3 4 5 6 Irradiation time (h)

Figure 3.30. Effect of solution pHs on lindane removal efficiency using S-TiO2

photocatalyst under visible light. [lindane]0 = 1.0 µM, [S-TiO2]0 = 0.23 g/L.

0.066 0.14 kobs

Initial degradation rate Initial degradation rate ( 0.064 0.12 0.062

) 0.10

-1 0.060

obs 0.058 0.08

0.056

0.06  0.054 M.h

-1 0.04 0.052 )

0.5 1.0 1.5 2.0

[Lindane]0 (M)

Figure 3.31. Variation of kobs and initial degradation rate with different initial concentration of lindane using S-TiO2 photocatalysis under visible light. The initial degradation rate corresponds to the first hour of decay. [S-TiO2]0 = 0.23 g/L, pH = 5.8.

184

1.0

0.8

0.6

(lindane) 0.6

)

-1 C/C 0.4 0.4

(h PMS/Vis 0.2

obs

k Ref-TiO2/PMS/Vis

0.2 0.0 S-TiO2/Vis S-TiO /PMS/Vis 0.0 0.2 0.4 0.6 0.8 1.0 2 [PMS] (mM) 0.0 0 1 2 3 4 5 6 Irradiation time (h)

Figure 3.32. Effect of 0.2 mM PMS on TiO2 photocatalysis of lindane under visible

light irradiation. [lindane]0 = 1.0 µM, pH = 5.8, [S-TiO2]0 = 0.23 g/L. Inset:

variation of kobs with [PMS]0 in the range of 0.1-1.0 mM.

Table 3.15. Pseudo first-order rate constant (kobs), removal efficiency (%) and half-life

(t1/2) of various visible light-assisted (VLA) AOPs based on S-TiO2 photocatalyst

for lindane degradation, in the presence and absence of 0.2 mM PMS. lindane]0 =

1.0 µM, pH = 5.8, [S-TiO2]0 = 0.23 g/L. ______

−1 Type of VLA-AOP kobs (h ) Removal efficiency (%) t1/2 (h)

−3 Ref-TiO2/vis 6.61 × 10 4.2 105.0

−2 Ref-TiO2/PMS/vis 1.19 × 10 7.3 58.2

−2 S-TiO2/vis 5.93 × 10 31.0 11.7

−1 S-TiO2/PMS/vis 1.87 × 10 68.2 3.7

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3.6.1.3 Effects of inorganic oxides on the degradation of lindane in the simulated solar/S-

TiO2 system

Photocatalytic degradation of lindane using S-TiO2 was also investigated under solar light irradiation and the results are shown in Figure 3.33. As seen in Figure 3.33, direct photodegradation of lindane by solar light irradiation was negligible within 6 hr.

However, the photocatalytic efficiency of the TiO2 based photocatalysts was significantly enhanced under solar irradiation, leading to 36.7 and 63.4% lindane removal using ref-

TiO2 and S-TiO2 films, respectively, for 6 hr. Solar radiation consists of about 5% UV light radiation, which has energy greater than the band gap energy of ref-TiO2, thereby capable of promoting electrons from valence band to the conduction band, generating an

− + electron-hole pair (eCB + hVB ) (reaction (4.75)). Subsequently, these electron-hole pairs

•− • can generate various ROSs such as O2 and OH, as discussed in the previous sections.

Although activation of ref-TiO2 by UV component of solar radiation resulted in a reasonable degree of lindane degradation, S-TiO2 photocatalysis yielded a far better result than the ref-TiO2 film mainly because of the potentially strong capacity of S-TiO2 film for absorbing visible light photons, besides, the increased BET surface area and porosity of the S-TiO2 film (Han et al., 2011). This is in accordance with the findings of Fotiou et al. (2015) and Triantis et al. (2012), who reported that various doped-TiO2 based photocatalysts showed higher performances than reference TiO2 for the degradation of pollutants under solar irradiation. The observed pseudo first-order rate constants for solar

−1 light-assisted ref-TiO2 and S-TiO2 photocatalysis were found to be 0.73 × 10 and 1.63 ×

10−1 h−1, respectively (Table 3.16).

186

The influence of PMS on solar light-assisted TiO2 photocatalysis of lindane was investigated (Figure 3.33), which showed a tendency similar to that of PMS effect on visible light-assisted S-TiO2 photocatalysis of lindan. Around 15.0% of lindane was decomposed by PMS (0.2 mM) direct oxidation under solar light for 6 hr, indicating that

PMS can be activated by solar light radiation. The photocatalytic activity of solar light- assisted TiO2 photocatalysis was significantly enhanced in the presence of PMS, with

85.4 and 99.9% lindane removal by ref-TiO2/PMS/solar and S-TiO2/PMS/solar processes, respectively, in 6 hour. The results shown here for the effect of PMS on solar light- assisted TiO2 photocatalysis of lindane are consistent with above discussion on the effect of PMS on visible light-assisted TiO2 photocatalysis of lindane (section 4.2.4), i.e.,

•− • reduced rate of the electron-hole recombination and generation of SO4 and/or OH.

Moreover, PMS can be directly activated by UV light (a fraction of solar radiation),

•− • thereby generating SO4 and OH radicals, as shown in reaction (4.90) (Anipsitakis and

−1 −1 −1 Dionysiou, 2004b). The value of kobs was found to be 3.10 × 10 and 5.83 × 10 h by ref-TiO2/PMS/solar and S-TiO2/PMS/solar processes, respectively (Table 3.16).

− •− • HSO5 + hν → SO4 + OH (4.90)

The effect of initial concentrations of PMS on the photocatalytic activity of S-

TiO2/solar process was also investigated, and the resulting kobs are shown in the inset of

Figure 3.33. The results showed that the value of kobs increased with increasing concentration of PMS, plausibly due to the reasons explained in the previous section, i.e.,

•− • increased SO4 and OH production as well as enhanced inhibition of electron-hole recombination at the increasing concentration of PMS. The calculated kobs was found to

−1 −1 −1 −1 −1 be 3.84 × 10 , 5.81 × 10 , 8.01 × 10 , and 9.97 × 10 h for S-TiO2/PMS/solar and

187

−1 −1 −1 −1 −1 that of 2.30 × 10 , 3.03 × 10 , 4.55 × 10 , and 5.71 × 10 h for ref-TiO2/PMS/solar when the initial concentration of PMS was 0.1, 0.2, 0.5 and 1.0 mM, respectively. An initial PMS concentration of 0.2 mM was used in all subsequent experiments, which yielded 85.4 and 99.9% lindane removal by ref-TiO2/PMS/solar and S-TiO2/PMS/solar respectively, in 6 hr.

A comparison of TiO2 photocatalysis in terms of percent removal efficiency (%) of lindane is illustrated in Table 3.16. The results of the present study showed that the removal efficiency of lindane followed the order: ref-TiO2/solar < S-TiO2/solar < ref-

TiO2/PMS/solar < S-TiO2/PMS/solar, which corresponded to 36.7, 63.4, 85.4 and 99.9% lindane removal, respectively in 6 hr. The observed pseudo first-order rate constant (kobs) and the calculated half-life (t1/2) for the processes studied herein are also given in Table

3.16. The results showed that solar light-assisted ref-TiO2 photocatalysis demonstrated high removal efficiency for lindane, which further increased in the presence of PMS, indicating that both ref-TiO2 and PMS were activated by solar light irradiation. The results also revealed that the half-life of the solar light-assisted TiO2 photocatalytic processes were even reduced in the presence of 0.2 mM PMS, thus making it more suitable for application purposes, by considerably reducing the size of the photocatalytic reactors.

188

1.0

0.8

PMS/solar Ref-TiO /PMS/solar 0.6 2 S-TiO2/PMS/solar

(lindane) 0 0.4 1.4 S-TiO2/PMS/solar C/C 1.2 Ref-TiO2/PMS/solar ) 1.0

-1 0.2 0.8

(h 0.6

obs

k 0.4 0.2 0.0 0.0 0.0 0.2 0.4 0.6 0.8 1.0 1.2 [PMS] (mM)

0 1 2 3 4 5 6 Irradiation time (h)

Figure 3.33. Solar light-assisted TiO2 and S-TiO2 photocatalysis of lindane in the

presence and absence of 0.2 mM PMS . [lindane]0 = 1.0 µM, pH = 5.8, [S-TiO2]0

= [TiO2]0 = 0.23 g/L. Inset: variation of kobs with [PMS]0 in the range of 0.1-1.0 mM.

Table 3.16. Pseudo first-order rate constant (kobs), removal efficiency (%) and half-life

(t1/2) of various solar light-assisted (SLA) AOPs based on S-TiO2 photocatalyst

for lindane degradation, in the presence and absence of 0.2 mM PMS. lindane]0 =

1.0 µM, pH = 5.8, [S-TiO2]0 = 0.23 g/L. ______

−1 Type of SLA-AOP kobs (h ) Removal efficiency (%) t1/2 (h)

−2 Ref-TiO2/solar 7.31 × 10 36.7 9.5

−1 S-TiO2/solar 1.63 × 10 63.4 4.3

−1 Ref-TiO2/PMS/solar 3.10 × 10 85.4 2.2

−1 S-TiO2/PMS/solar 5.83 × 10 99.9 1.2

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4 CONCLUSIONS AND FUTURE PERSPECTIVES

4.3 Conclusions

The effective degradation of lindane by gamma irradiation was plausibly due to its

− 10 ˗1 ˗1 high second-order rate constant with eaq , i.e., 1.26 ± 0.04 × 10 M s , as determined by pulse radiolysis techniques. The concomitant background constituents, such as NOM and various inorganic ions showed significant influences on the process efficiency, indicating a need for pretreatment, aiming at the removal of those constituents for the purpose of practical applications. The pH of the solution also had a considerable effect on the process efficiency. The highest removal efficiency was obtained at neutral pH (6.8).

Despite high inhibition from the concomitant background constituents, effective lindane removal could still be achieved, though at relatively higher radiation doses. The nearly complete dechlorination of lindane may indicate the gamma radiation based AO/RPs to be a promising technology for water detoxification, considering the common relationship of chlorinated organic compounds with toxicity.

•− The SO4 radical reacted with lindane with a second-order rate constant of 1.3 ×

109 M−1 s−1, as determined by competition kinetics, making the degradation of lindane to be plausible by UV/PMS based AOPs. The observed pseudo-first order rate constant

(kobs) was affected by the initial concentrations of PMS and lindane as well as the solution

2− − pH. The presence of humic acid, CO3 or HCO3 caused a strong inhibiting effect while

2− − the presence of SO4 or Cl exerted a negligible effect on the efficiency of UV/PMS process. Various degradation by-products, as identified by GC/MS, revealed dechlorination, chlorination, dehydrogenation and hydroxylation to be potential transformation pathways. Ring opening and cleavage could also be achieved indirectly as 190

demonstrated by the significant decrease in TOC. This study shows UV/PMS based AOP is an effective method for the removal of lindane from aqueous solution.

2+ − The Fe /HSO5 based processes was evaluated for the degradation and mineralization of lindane from aqueous solution and found to be enhanced in the presence of fluorescence light and also by UV light irradiation. The lindane degradation followed pseudo-first order kinetics in the tested AOPs. The efficiency of the

2+ − 2+ − UV/Fe /HSO5 system increased with increasing concentrations of either Fe or HSO5 but decreased with increasing pH. The observed pseudo-first order rate constant of

2+ − lindane by UV/Fe /HSO5 was found to decrease with increasing concentration of contaminant while the initial degradation rate increased with higher lindane

2+ − 2+ − concentrations. The Fe /HSO5 and fluorescence light/Fe /HSO5 processes showed

2+ − limited ability in the mineralization of lindane, while UV/Fe /HSO5 system presented a

− much better TOC removal efficiency. HSO5 was readily decomposed, especially in the systems containing both Fe2+ and UV radiation, consistent with lindane degradation and

2+ − mineralization efficiencies. Overall, considering degradation kinetics, UV/Fe /HSO5 system is the most efficient for the degradation and mineralization of lindane in this study.

Nanocrystalline S-TiO2 films, synthesized by a sol–gel method, exhibited significant photocatalytic efficiency for the degradation of lindane under visible light irradiation. The efficiency of S-TiO2 photocatalysis was affected by the concentration of the contaminant and the solution pH. The highest photocatalytic degradation of lindane was observed at near neutral pH (5.8). TiO2 photocatalysis significantly improved under

− solar irradiation. The addition of 0.2 mM HSO5 showed a remarkable enhancing effect

191

on TiO2 photocatalysis of lindane, causing many fold reduction in the half-lives of the

− reaction. Therefore, by adding HSO5 , sizes of photocatalytic reactors can be considerably reduced in practical applications. Based on the results, key operating parameters were optimized. The results indicated that solar or visible light-active TiO2- based AOPs are very effective for the detoxification of water contaminated with chlorinated pesticides, such as lindane.

The surface water samples collected from district Swabi, Khyber Pakhtoonkhwa,

Pakistan, showed high concentration of lindane. The high lindane residues in the field water samples of district Swabi shows that despite ban, lindane is still being used in

Pakistan for agriculture purposes. Canal water was comparatively highly contaminated with lindane, probably due to the discharges from the agriculture fields.

4.4 Future Perspectives

The relatively positive influence of fluorescence light (e.g., tube-light) on the studied

Fenton-like process may suggest a further research need and a potential application of

2+ − Fe /HSO5 in field, by using other long wavelength light sources such as UV-vis

(including solar light) and UVA. The photocatalytic activity of the synthesized S-TiO2 may be evaluated for the removal of other pollutants, particularly the chlorinated organic compounds. The effect of natural sun light on photocatalytic activity of S-TiO2 for lindane degradation may also be investigated. The efficiency of the studied AOPs may also be checked for the treatment of the field waters contaminated with lindane, which probably will provide useful scientific informations on the extension of AOPs in the real- world applications. The relatively high lindane concentrations detected in the analyzed water samples may suggest proper education should be given to the farmers about the 192

hazardous effects of this toxic pesticide to avoid its use in future. The residence of the region should be warned about the use of surface water for drinking purposes. More detailed study is needed to explore such other toxic pesticides prevailing in this region.

The current study can be extended to other regions of Pakistan in order to get a clear baseline data for the entire country. Based on the results obtained, it is recommended that the government of Pakistan should take strong action against the use of these persistent toxic pesticides for agriculture purposes.

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REFRENCES

Abhilash, P.C., Singh, N., 2009. Pesticide use and application: An Indian scenario. J. Hazard. Mater. 165, 1-12. Agrawal A., Pandey R. S, Sharma B., 2010. Water Pollution with Special Reference to Pesticide Contamination in India. J. Water Res.our Prot. 02, 432-448. Agrawal, A., Tratnyek, P.G., 1995. Reduction of Nitro Aromatic Compounds by Zero- Valent Iron Metal. Environ. Sci. Technol. 30, 153-160. Akpan, U.G., Hameed, B.H., 2009. Parameters affecting the photocatalytic

degradation of dyes using TiO2-based photocatalysts: A review. Journal of

hazardous materials 170, 520-529.

Al-Chalabi, K.A.K., Al-Khayat, B.H.A., 1989. The effect of lindane on nucleic acids, protein and carbohydrate content in Tetrahymena pyriformis. Environ. Pollut. 57, 281-287. Alegre M.L., Geronés, M., Rosso, J.A., Bertolotti, S.G., Braun, A.M., Mártire,

D.O., Gonzalez, M.C., 2000. Kinetic Study of the Reactions of Chlorine Atoms

and Cl2•- Radical Anions in Aqueous Solutions. 1. Reaction with Benzene. The

Journal of Physical Chemistry A 104, 3117-3125.

Alexander, M., 1994. Biodegradataion bioremediation: 2nd ed. Academic Press. Alinsafi, A., Khemis, M., Pons, M.N., Leclerc, J.P., Yaacoubi, A., Benhammou, A., Nejmeddine, A., 2005. Electro-coagulation of reactive textile dyes and textile wastewater. Chem. Eng. Process: Process Intensif. 44, 461-470. Andersen, J., Pelaez, M., Guay, L., Zhang, Z., O'Shea, K., Dionysiou, D.D., 2013. NF-

TiO2 photocatalysis of amitrole and atrazine with addition of oxidants under simulated solar light: emerging synergies, degradation intermediates, and reusable attributes. J. Hazard. Mater.260, 569-575. Angelini, G., Bucci, R., Carnevaletti, F., Colosimo, M., 2000. Radiolytic decomposition of aqueous atrazine. Radiat. Phys. Chem. 59, 303-307.

194

Anipsitakis, G.P., Dionysiou, D.D., 2003. Degradation of organic contaminants in water with sulfate radicals generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci. Technol. 37, 4790-4797. Anipsitakis, G.P., Dionysiou, D.D., 2004a. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 38, 3705-3712. Anipsitakis, G.P., Dionysiou, D.D., 2004b. Transition metal/UV-based advanced oxidation technologies for water decontamination. Appl. Catal. B-Environ 54, 155-163. Anipsitakis, G.P., Stathatos, E., Dionysiou, D.D., 2005. Heterogeneous activation of

Oxone using Co3O4. J. Phys. Chem. B 109, 13052-13055. Anipsitakis, G.P., Dionysiou, D.D., Gonzalez, M.A., 2006. Cobalt-mediated activation of peroxymonosulfate and sulfate radical attack on phenolic compounds. implications of chloride ions. Environ. Sci. Technol. 40, 1000-1007. Antonaraki, S., Triantis, T.M., Papaconstantinou, E., Hiskia, A., 2010. Photocatalytic degradation of lindane by polyoxometalates: Intermediates and mechanistic aspects. Catal. Today 151, 119-124. Antoniou, M.G., de la Cruz, A.A., Dionysiou, D.D., 2010a. Degradation of microcystin- LR using sulfate radicals generated through photolysis, thermolysis and e− transfer mechanisms. Appl. Catal. B: Environ. 96, 290-298. Antoniou, M.G., de la Cruz, A.A., Dionysiou, D.D., 2010b. Intermediates and Reaction pathways from the Degradation of Microcystin-LR with Sulfate Radicals. Environ. Sci. Technol. 44, 7238-7244. Asahi, R., Morikawa, T., Ohwaki, T., Aoki, K., Taga, Y., 2001. Visible-light photocatalysis in Nitrogen-doped Titanium oxides. Science 293, 269-271. ATSDR (2005) Toxicological profile for alpha-, beta-, gamma- and delta hexachlorocyclohexane. Atlanta, GA. US Department of Health and Human Services, Public Health Service, Agency for Toxic Substances and Disease Registry. Azizullah, A., Khattak, M.N., Richter, P., Hader, D.P., 2011. Water pollution in Pakistan and its impact on public health--a review. Environ. Int. 37, 479-497.

195

Balkaş, T.I., 1972. The radiolysis of aqueous solutions of methylene chloride. Int. J. Radiat. Phys. Chem. 4, 199-208. Bandala, E.R., Pelaez, M.A., Salgado, M.J., Torres, L., 2008. Degradation of sodium dodecyl sulphate in water using solar driven Fenton-like advanced oxidation processes. J. Hazard. Mater.151, 578-584. Bandala, E.R., Brito, L., Pelaez, M., 2009. Degradation of domoic acid toxin by UV- promoted Fenton-like processes in seawater. Desalination 245, 135-145. Banerjee, S., Pillai, S.C., Falaras, P., O’shea, K.E., Byrne, J.A., Dionysiou, D.D.,

2014. New insights into the mechanism of visible light photocatalysis. The

Journal of Physical Chemistry Letters 5, 2543-2554.

Barber, J.L., Sweetman, A.J., van Wijk, D., Jones, K.C., 2005. Hexachlorobenzene in the global environment: Emissions, levels, distribution, trends and processes. Sci. Total Environ. 349, 1-44. Barrie, L.A., Gregor, D., Hargrave, B., Lake, R., Muir, D., Shearer, R., Tracey, B., Bidleman, T., 1992. Arctic contaminants: sources, occurrence and pathways. Sci. Total Environ. 122, 1-74. Basfar, A.A., Khan, H.M., Al-Shahrani, A.A., 2005a. Trihalomethane treatment using gamma irradiation: kinetic modeling of single solute and mixtures. Radiat. Phys. Chem. 72, 555-563. Basfar, A.A., Khan, H.M., Al-Shahrani, A.A., Cooper, W.J., 2005b. Radiation induced decomposition of methyl tert-butyl ether in water in presence of chloroform: Kinetic modelling. Water Res.. 39, 2085-2095. Basfar, A.A., Mohamed, K.A., Al-Abduly, A.J., Al-Kuraiji, T.S., Al-Shahrani, A.A., 2007. Degradation of diazinon contaminated waters by ionizing radiation. Radiat. Phys. Chem. 76, 1474-1479. Bauer, R., Fallmann, H., 1997. The Photo-Fenton Oxidation — A cheap and efficient wastewater treatment method. Res. Chem. Intermed. 23, 341-354. Behnajady, M.A., Modirshahla, N., Shokri, M., 2004. Photodestruction of Acid Orange 7

196

(AO7) in aqueous solutions by UV/H2O2: influence of operational parameters. Chemosphere 55, 129-134. Beland, F.A., Farwell, S.O., Callis, P.R., Geer, R.D., 1977. Reduction pathways of organohalogen compounds: Part III. A molecular orbital (CNDO/2) study of the chlorinated benzenes, DDT, and lindane. J. Electroanalyt. Chem. Interfac. Electrochem. 78, 145-159. Beland, F.A., Farwell, S.O., Robocker, A.E., Geer, R.D., 1976. Electrochemical reduction and anaerobic degradation of lindane. J. Agric. Food Chem.. 24, 753- 756. Begum, A., Gautam, S.K., 2012. Endosulfan and lindane degradation using ozonation. Environ. Technol. 33, 943-949. Beltrán, F.J., Encinar, J.M., Alonso, M.A., 1998. Nitroaromatic hydrocarbon ozonation in water. 2. Combined Ozonation with Hydrogen Peroxide or UV Radiation. Indust. Engin. Chem. Res. 37, 32-40. Benitez, F.J., Real, F.J., Acero, J.L., Garcia, C., 2006. Photochemical oxidation processes for the elimination of phenyl-urea herbicides in waters. J. Hazard. Mater. 138, 278-287. Benimeli, C.S., Fuentes, M.S., Abate, C.M., Amoroso, M.J., 2008. Bioremediation of lindane-contaminated soil by Streptomyces sp M7 and its effects on Zea mays growth. Int. Biodeterior. Biodegrad. 61, 233-239.

Bertelli, M., Selli, E., 2004. Kinetic analysis on the combined use of photocatalysis, H2O2 photolysis, and sonolysis in the degradation of methyl tert-butyl ether. Appl. Catal. B: Environ. 52, 205-212. Betterton, E.A., Hoffmann, M.R., 1990. Kinetics and Mechanism of the Oxidation of Aqueous hydrogen-sulfide by Peroxymonosulfate. Environ. Sci. Technol. 24, 1819-1824. Bhatkhande, D.S., Pangarkar, V.G., Beenackers, A.A.C.M., 2002. Photocatalytic

degradation for environmental applications – a review. Journal of Chemical

Technology & Biotechnology 77, 102-116.

197

Birke, V., Mattik, J., Runne, D., 2004. Mechanochemical reductive dehalogenation of hazardous polyhalogenated contaminants. J. Mater. Sci. 39, 5111-5116. Bojanowska-Czajka, A., Drzewicz, P., Kozyra, C., Nalecz-Jawecki, G., Sawicki, J., Szostek, B., Trojanowicz, M., 2006. Radiolytic degradation of herbicide 4-chloro- 2-methyl phenoxyacetic acid (MCPA) by gamma-radiation for environmental protection. Ecotoxicol. Environ. Saf. 65, 265-277. Bojanowska-Czajka, A., Drzewicz, P., Zimek, Z., Nichipor, H., Nałęcz-Jawecki, G., Sawicki, J., Kozyra, C., Trojanowicz, M., 2007. Radiolytic degradation of pesticide 4-chloro-2-methylphenoxyacetic acid (MCPA)—Experimental data and kinetic modelling. Radiat. Phys. Chem. 76, 1806-1814. Bolton, J.R., 2011. Ultraviolet Applications Handbook. Bolton Photosciences, 3rd ed. Edmonton, Alta. Bossmann, S.H., Oliveros, E., Göb, S., Siegwart, S., Dahlen, E.P., Payawan, L., Straub, M., Wörner, M., Braun, A.M., 1998. New evidence against hydroxyl radicals as reactive intermediates in the thermal and photochemically enhanced Fenton reactions. J. Phys. Chem. A 102, 5542-5550. Brinkmann, T., Horsch, P., Sartorius, D., Frimmel, F.H., 2003. Photoformation of low- molecular-weight organic acids from brown water dissolved organic matter. Environ. Sci. Technol. 37, 4190-4198. Burbano, A.A., Dionysiou, D.D., Suidan, M.T., 2008. Effect of oxidant-to-substrate ratios on the degradation of MTBE with Fenton reagent. Water Res. 42, 3225- 3239. Buxton, G.V., Greenstock, C.L., Helman, W.P., Ross, A.B., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (•OH/•O−) in aqueous solution. J. Phys. Chem. Ref. Data 17, 513-886. Camel, V., Bermond, A., 1998. The use of ozone and associated oxidation processes in drinking water treatment. Water Res. 32, 3208-3222. Cao, J.S., Zhang, W.X., Brown, D.G., Sethi, D., 2008. Oxidation of lindane with Fe(II)- activated sodium persulfate. Environ. Eng. Sci. 25, 221-228. Carey, J.H., Lawrence, J., Tosine, H.M., 1976. Photodechlorination of PCB's in the

198

presence of titanium dioxide in aqueous suspensions. Bull. Environ. Contam. Toxicol. 16, 697-701. Cha, J.-A., An, S.-H., Jang, H.-D., Kim, C.-S., Song, D.-K., Kim, T.-O., 2012. Synthesis

and photocatalytic activity of N-doped TiO2/ZrO2 visible-light photocatalysts. Adv. Powder Technol. 23, 717-723. Chan, K.H., Chu, W., 2009. Degradation of atrazine by cobalt-mediated activation of peroxymonosulfate: Different cobalt counteranions in homogenous process and cobalt oxide catalysts in photolytic heterogeneous process. Water Res. 43, 2513- 2521. Chatterjee, D., Mahata, A., 2004. Evidence of superoxide radical formation in the

photodegradation of pesticide on the dye modified TiO2 surface using visible light. J. Photochem. Photobiol. A: Chem. 165, 19-23. Chelme-Ayala, P., El-Din, M.G., Smith, D.W., 2010. Degradation of bromoxynil and

trifluralin in natural water by direct photolysis and UV plus H2O2 advanced oxidation process. Water Res. 44, 2221-2228. Chen, X., Qiao, X., Wang, D., Lin, J., Chen, J., 2007. Kinetics of oxidative decolorization and mineralization of Acid Orange 7 by dark and photoassisted Co2+-catalyzed peroxymonosulfate system. Chemosphere 67, 802-808. Chen, X., Wang, W., Xiao, H., Hong, C., Zhu, F., Yao, Y., Xue, Z., 2012. Accelerated

TiO2 photocatalytic degradation of Acid Orange 7 under visible light mediated by peroxymonosulfate. Chem. Eng. J. 193–194, 290-295. Chen, H., Bramanti, E., Longo, I., Onor, M., Ferrari, C., 2011. Oxidative decomposition of atrazine in water in the presence of hydrogen peroxide using an innovative microwave photochemical reactor. J. Hazard. Mater. 186, 1808-1815. Cheng, X., Yu, X., Xing, Z., Yang, L., 2012. Synthesis and characterization of N-doped

TiO2 and its enhanced visible-light photocatalytic activity. Arab. j. Chem. xxx, xxx-xxx. Chevaldonnet, C., Cardy, H., Dargelos, A., 1986. Ab initio CI calculations on the PE and VUV spectra of hydrogen peroxide. Chem. Phys. 102, 55-61. Chian, E.S.K., Bruce, W.N., Fang, H.H.P., 1975. Removal of Pesticides by Reverse-

199

Osmosis. Environ. Sci. Technol. 9, 52-59. Choi, D., Lee, O.M., Yu, S., Jeong, S.W., 2010. Gamma radiolysis of alachlor aqueous solutions in the presence of hydrogen peroxide. J. Hazard. Mater. 184, 308-312. Chong, M.N., Jin, B., Chow, C.W., Saint, C., 2010. Recent developments in photocatalytic water treatment technology: a review. Water Res. 44, 2997-3027. Christensen, H., Sehested, K., Corfitzen, H., 1982. Reactions of hydroxyl radicals with hydrogen peroxide at ambient and elevated temperatures. J. Phys. Chem. 86, 1588-1590. Chu, L.B., Wang, J.L., Wang, B., 2010. Effects of aeration on gamma irradiation of sewage sludge. Radiat. Phys. Chem. 79, 912-914. Chu, W., Wang, Y.R., Leung, H.F., 2011. Synergy of sulfate and hydroxyl radicals in 2− UV/S2O8 /H2O2 oxidation of iodinated X-ray contrast medium iopromide. Chem. Eng. J. 178, 154-160. Comerton, A.M., Andrews, R.C., Bagley, D.M., 2005. Evaluation of an MBR–RO system to produce high quality reuse water: Microbial control, DBP formation and nitrate. Water Res. 39, 3982-3990. Comninellis, C., 1994. Electrocatalysis in the electrochemical conversion/combustion of organic pollutants for waste water treatment. Electrochim. Acta 39, 1857-1862. Cooper, W.J., Mezyk, S.P., O’Shea, K.E., Kim, D.K., Hardison, D.R., 2003. Kinetic modeling of the destruction of methyl tert-butyl ether (MTBE). Radiat. Phys. Chem. 67, 523-526. Crawford, R.L., Crawford, D.L., 2005. Bioremediation: Principles and Applications. Cambridge University Press. Criquet, J., Karpel Vel Leitner, N., 2011. Radiolysis of acetic acid aqueous solutions— Effect of pH and persulfate addition. Chem. Eng. J. 174, 504-509. Criquet, J., Leitner, N.K., 2009. Degradation of acetic acid with sulfate radical generated by persulfate ions photolysis. Chemosphere 77, 194-200. Cristol, S.J., 1947. The kinetics of the alkaline dehydrochlorination of the benzene hexachloride isomers. The mechanism of second-order elimination reactions1,2. J. Am. Chem. Soc. 69, 338-342.

200

Criswell, K.A., Loch–Caruso, R., 1999. Lindane-induced inhibition of spontaneous contractions of pregnant rat uterus. Reproduct. Toxicol. 13, 481-490. Csay, T., Rácz, G., Takács, E., Wojnárovits, L., 2012. Radiation induced degradation of pharmaceutical residues in water: Chloramphenicol. Radiat. Phys. Chem. 81, 1489-1494. Daneshvar, N., Behnajady, M.A., Zorriyeh Asghar, Y., 2007. Photooxidative degradation

of 4-nitrophenol (4-NP) in UV/H2O2 process: Influence of operational parameters and reaction mechanism. J. Hazard. Mater. 139, 275-279. Dantas, R.F., Rossiter, O., Teixeira, A.K.R., Simões, A.S.M., da Silva, V.L., 2010. Direct UV photolysis of propranolol and metronidazole in aqueous solution. Chem. Eng. J. 158, 143-147. De Gusseme, B., Soetaert, M., Hennebel, T., Vanhaecke, L., Boon, N., Verstraete, W., 2012. Catalytic dechlorination of diclofenac by biogenic palladium in a microbial electrolysis cell. Microb. biotechnol. 5, 396-402. Devi, L.G., Kavitha, R., 2013. A review on non metal ion doped Titania for the photocatalytic degradation of organic pollutants under UV/solar light: Role of photogenerated charge carrier dynamics in enhancing the activity. Appl. Catal. B: Environ.140–141, 559-587. Di Valentin, C., Finazzi, E., Pacchioni, G., Selloni, A., Livraghi, S., Paganini, M.C.,

Giamello, E., 2007. N-doped TiO2: Theory and experiment. Chem. Phys.339, 44- 56. Dionysiou, D., Khodadoust, A.P., Kern, A.M., Suidan, M.T., Baudin, I., Laine, J.M., 2000. Continuous-mode photocatalytic degradation of chlorinated phenols and

pesticides in water using a bench-scale TiO2 rotating disk reactor. Appl. Catal. B- Environ 24, 139-155. Donald, D.B., Block, H., Wood, J., 1997. Role of ground water on hexachlorocyclohexane (lindane) detections in surface water in western Canada. Environ. Toxicol. Chem. 16, 1867-1872. Drijvers, D., Van Langenhove, H., Vervaet, K., 1998. Sonolysis of chlorobenzene

in aqueous solution: organic intermediates. Ultrasonics Sonochemistry 5, 13-19.

201

Drzewicz, P., Nalecz-Jawecki, G., Gryz, M., Sawicki, J., Bojanowska-Czajka, A., Gluszewski, W., Kulisa, K., Wolkowicz, S., Trojanowicz, M., 2004. Monitoring of toxicity during degradation of selected pesticides using ionizing radiation. Chemosphere 57, 135-145. El-Dib, M.A., Ramadan, F.M., Ismail, M., 1975. Adsorption of Sevin and Baygon on granular activated carbon. Water Res. 9, 795-798.

Elkanzi, E.M., Bee Kheng, G., 2000. H2O2/UV degradation kinetics of isoprene in aqueous solution. J. Hazard. Mater. 73, 55-62. Elmolla, E., Chaudhuri, M., 2009. Optimization of Fenton process for treatment of amoxicillin, ampicillin and cloxacillin antibiotics in aqueous solution. J. Hazard. Mater.170, 666-672. Fan, X., Hao, H., Shen, X., Chen, F., Zhang, J., 2011. Removal and degradation pathway study of sulfasalazine with Fenton-like reaction. J. Hazard. Mater.190, 493-500. Faust, B.C., Hoigné, J., 1990. Photolysis of Fe (III)-hydroxy complexes as sources of •OH radicals in clouds, fog and rain. Atmos. Environ. A: Gen. Top. 24, 79-89. Fernandez, J., Maruthamuthu, P., Kiwi, J., 2004. Photobleaching and mineralization of Orange II by oxone and metal-ions involving Fenton-like chemistry under visible light. J. Photoch. Photobio. A. 161, 185-192. Fenton, H.J.H., 1894. Oxidation of tartaric acid in presence of iron. J. Chem. Soc., Trans. 65, 899-910. Flanagan, J., Griffith, W.P., Skapski, A.C., 1984. The active principle of Caro's acid, − HSO5 : X-ray crystal structure of KHSO5.H2O. J. Chem. Soc., Chem. Commun., 1574-1575. Fotiou, T., Triantis, T.M., Kaloudis, T., Hiskia, A., 2015. Evaluation of the

photocatalytic activity of TiO2 based catalysts for the degradation and

mineralization of cyanobacterial toxins and water off-odor compounds under UV-

A, solar and visible light. Chemical Engineering Journal 261, 17-26.

Fox, M.A., Dulay, M.T., 1993. Heterogeneous photocatalysis. Chem. Rev. 93, 341-357. Gao, Y.Q., Gao, N.Y., Deng, Y., Yang, Y.Q., Ma, Y., 2012. Ultraviolet (UV) light-

202

activated persulfate oxidation of sulfamethazine in water. Chem. Eng. J. 195, 248- 253. Gehringer, P., Proksch, E., Eschweiler, H., Szinovatz, W., 1990. Removal of chlorinated ethylenes from drinking water by radiation treatments. Int. J. Rad. Appl. and Instrum. C. Radiat. Phys. Chem. 35, 456-460. Getoff, N., 1986. Radiation induced decomposition of some chlorinated methanes in water. Water Res. 20, 1261-1264. Getoff, N., 1995. Radiation-induced degradation of water pollutants: State of the art. Radiat. Phys. Chem. 46, 1079. Getoff, N., Lutz, W., 1985. Radiation induced decomposition of hydrocarbons in water resources. Radiat. Phys. Chem. 25 (1), 21–26. Ghaly, M.Y., Hartel, G., Mayer, R., Haseneder, R., 2001. Photochemical oxidation of p-

chlorophenol by UV/H2O2 and photo-Fenton process. A comparative study. Waste Manag. 21, 41-47. Ghodbane, H., Hamdaoui, O., 2010. Decolorization of antraquinonic dye, C.I.

Acid Blue 25, in aqueous solution by direct UV irradiation, UV/H2O2 and

UV/Fe(II) processes. Chemical Engineering Journal 160, 226-231.

Gilbert, B.C., Stell, J.K., Peet, W.J., Radford, K.J., 1988. Generation and

reactions of the chlorine atom in aqueous solution. Journal of the Chemical

Society, Faraday Transactions 1: Physical Chemistry in Condensed Phases 84,

3319-3330.

Glaze, W.H., 1987. Drinking-water treatment with ozone. Environ. Sci. Technol. 21, 224- 230. Gogate, P.R., Pandit, A.B., 2004. A review of imperative technologies for wastewater treatment II: hybrid methods. Adv. Environ. Res. 8, 553-597. Goldstein, S., Rabani, J., 2008. The ferrioxalate and iodide–iodate actinometers in the UV region. J. Photochem. Photobiol. A: Chem.193, 50-55. Guan, Y.H., Ma, J., Ren, Y.M., Liu, Y.L., Xiao, J.Y., Lin, L.Q., Zhang, C., 2013.

203

Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals. Water Res. 47, 5431-5438. Georgiou, D., Melidis, P., Aivasidis, A., 2003. Use of a microbial sensor: inhibition effect of azo-reactive dyes on activated sludge, Bioprocess Biosyst. Eng. 25, 79– 83. Guan, Y.H., Ma, J., Ren, Y.M., Liu, Y.L., Xiao, J.Y., Lin, L.Q., Zhang, C., 2013. Efficient degradation of atrazine by magnetic porous copper ferrite catalyzed peroxymonosulfate oxidation via the formation of hydroxyl and sulfate radicals. Water Res. 47, 5431-5438. Guillard, C, Huber, P.P., G, Hoang-Van, C, 1996. The GC-MS analysis of organic

intermediates from the TiO2 photocatalytic treatment of water contaminated by

lindane (1a, 2a, 3b, 4a, 5a,6b-hexachlorocyclohexane). J. Adv. Oxid. Technol. 1,

53-60.

Guo, Y., Wang, Y., Hu, C., Wang, Y., Wang, E., Zhou, Y., Feng, S., 2000. Microporous

Polyoxometalates POMs/SiO2: Synthesis and photocatalytic degradation of aqueous organocholorine pesticides. Chem. Mater. 12, 3501-3508. Guo, Z., Zheng, Z., Gu, C., Tang, D., 2009. Radiation removals of low-concentration halomethanes in drinking water. J. Hazard. Mater.164, 900-903. Gupta, A., Kaushik, C.P., Kaushik, A., 2000. Degradation of hexachlorocyclohexane (HCH; α, β, γ and δ) by Bacillus circulans and Bacillus brevis isolated from soil contaminated with HCH. Soil Biol. Biochem. 32, 1803-1805. Gupta, K.C., Sutar, A.K., 2008. Polymer supported catalysts for oxidation of phenol and cyclohexene using hydrogen peroxide as oxidant. J Mol. Catal. A: Chem 280, 173-185. Güsten, H., Filby, W.G., Schoof, S., 1981. Prediction of hydroxyl radical reaction rates with organic compounds in the gas phase. Atmos. Environ. (1967) 15, 1763-1765. Haag, W.R., Yao, C.C.D., 1992. Rate constants for reaction of hydroxyl radicals with several drinking-water contaminants. Environ. Sci. Technol.26, 1005-1013.

204

Hall, A.K., Harrowfield, J.M., Hart, R.J., McCormick, P.G., 1996. Mechanochemical reaction of DDT with calcium oxide. Environ. Sci. Technol.30, 3401-3407. Han, C., Pelaez, M., Likodimos, V., Kontos, A.G., Falaras, P., O'Shea, K., Dionysiou,

D.D., 2011. Innovative visible light-activated sulfur doped TiO2 films for water treatment. Appl. Catal. B: Environ.107, 77-87. Hansen, L.G., 1998. Stepping backward to improve assessment of PCB congener toxicities. Environ. Health Perspect. 106, 171-189 Hargrave, B.T., Vass, W.P., Erickson, P.E., Fowler, B.R., 1988. Atmospheric transport of organochlorines to the Arctic Ocean. Tellus B 40B, 480-493. Hashmi, I., Farooq, S., Qaiser, S., 2009. Chlorination and water quality monitoring within a public drinking water supply in Rawalpindi Cantt (Westridge and Tench) area, Pakistan. Environ. Monit. Assess. 158, 393-403. Hayes, W.J., 1982. Pesticides studied in man. Williams and Wilkins. Hernandez, R., Zappi, M., Kuo, C.-H., 2004. Chloride effect on TNT degradation by Zero-valent Iron or Zinc during water treatment. Environ. Sci. Technol.38, 5157- 5163. Hessler, D.P., Gorenflo, V., Frimmel, F.H., 1993. Degradation of aqueous Atrazine and

Metazachlor solutions by UV and UV/H2O2 — Influence of pH and herbicide concentration. Acta Hydroch. Hydrob. 21, 209-214. Hijnen, W.A., Beerendonk, E.F., Medema, G.J., 2006. Inactivation credit of UV radiation for viruses, bacteria and protozoan (oo)cysts in water: a review. Water Res. 40, 3- 22. Hill, D.W., McCarty, P.L., 1967. Anaerobic degradation of selected chlorinated hydrocarbon pesticides. J. Water Pollut. Control Federation 39, 1259-1277. Hinrichsen, D., Tacio, H., 1997. Environmental Change and Security Program. The coming freshwater crisis is already here. 1-26, Wilson Center, Washington DC. Hiskia, A., Androulaki, E., Mylonas, A., Boyatzis, S., Dimoticali, D., Minero, C.,

Pelizzetti, E., Papaconstantinou, E., 2000. Photocatalytic mineralization of

chlorinated organic pollutants in water by polyoxometallates. Determination of

205

intermediates and final degradation products. Research on Chemical

Intermediates 26, 235-251.

Hislop, K.A., Bolton, J.R., 1999. The photochemical generation of hydroxyl radicals in

the UV−vis/Ferrioxalate/H2O2 system. Environ. Sci. Technol.33, 3119-3126. Ho, P.C., 1986. Photooxidation of 2,4-dinitrotoluene in aqueous solution in the presence of hydrogen peroxide. Environ. Sci. Technol.20, 260-267. Hoffmann, M.R., Martin, S.T., Choi, W., Bahnemann, D.W., 1995. Environmental applications of semiconductor photocatalysis. Chem. Rev. 95, 69-96. Hooftman, R.N., de Kreuk J.F., 1982. Investigation of the environmental load of chlorinated benzenes (Literature study). TNO, Netherlands. Hoigné, J., 1998. Chemistry of Aqueous Ozone and Transformation of Pollutants by Ozonation and Advanced Oxidation Processes. in: Hrubec, J. (Ed.). Quality and Treatment of Drinking Water II. Springer Berlin Heidelberg, pp. 83-141. Holm, N.W., Berry, R.J., 1970. Manual on radiation dosimetry. M. Dekker. Hosomi, M., 2001. Adoption and future point at issue of the Stockholm convention on persistent organic pollutants. Waste Manage. Res. 12, 338–348. Hosomi, M., 2002. Adoption of the Stockholm convention on persistent organic pollutants. J. Water Waste 44, 219–225. Huang, K.-C., Couttenye, R.A., Hoag, G.E., 2002. Kinetics of heat-assisted persulfate oxidation of methyl tert-butyl ether (MTBE). Chemosphere 49, 413-420. Huang, Y.F., Huang, Y.H., 2009. Behavioral evidence of the dominant radicals and intermediates involved in bisphenol A degradation using an efficient Co2+/PMS oxidation process. J. Hazard. Mater.167, 418-426. Huang, Y.-F., Huang, Y.-H., 2009. Identification of produced powerful radicals involved

in the mineralization of bisphenol A using a novel UV-Na2S2O8/H2O2-Fe(II,III) two-stage oxidation process. J. Hazard. Mater.162, 1211-1216. Huang, Y.H., Huang, Y.F., Huang, C.I., Chen, C.Y., 2009. Efficient decolorization of azo dye Reactive Black B involving aromatic fragment degradation in buffered Co2+/PMS oxidative processes with a ppb level dosage of Co2+-catalyst. J. Hazard. Mater.170, 1110-1118. 206

Huber, A., Bach, M., Frede, H.G., 2000. Pollution of surface waters with pesticides in Germany: modeling non-point source inputs. Agric. Ecosyst. Environ. 80, 191- 204. Huie, R.E., Clifton, C.L., Neta, P., 1991. Electron-transfer reaction-rates and equilibria of the carbonate and sulfate radical-anions. Radiat. Phys. Chem. 38, 477-481. Idaka, E., Ogawa, T., Horitsu, H., 1987. Reductive metabolism of aminoazobenzenes By Pseudomonas cepacia. Bull. Environ. Contam. Toxicol. 39, 100-107. Iwata, H., Tanabe, S., Sakai, N., Nishimura, A., Tatsukawa, R., 1994. Geographical distribution of persistent organochlorines in air, water and sediments from Asia and Oceania, and their implications for global redistribution from lower latitudes. Environ. Pollut. 85, 15-33. Jagnow, G., Haider, K., Ellwardt, P.C., 1977. Anaerobic dechlorination and degradation of hexachlorocyclohexane isomers by anaerobic and facultative anaerobic bacteria. Arch. Microbiol. 115, 285-292. Javier Benitez, F., Acero, J.L., Real, F.J., 2002. Degradation of carbofuran by using ozone, UV radiation and advanced oxidation processes. J. Hazard. Mater. 89, 51- 65. Jay, K., Stieglitz, L., 1995. Identification and quantification of volatile organic components in emissions of waste incineration plants. Chemosphere 30, 1249– 1260. Johri, A.K., Dua, M., Tuteja, D., Saxena, R., Saxena, D.M., Lal, R., 1996. Genetic manipulations of microorganisms for the degradation of hexachlorocyclohexane. FEMS microbial. Rev. 19, 69-84. Jones, K.C., de Voogt, P., 1999. Persistent organic pollutants (POPs): state of the science. Environ. Pollut. 100, 209-221. Kalsoom, U., Ashraf, S.S., Meetani, M.A., Rauf, M.A., Bhatti, H.N., 2012. Degradation

and kinetics of H2O2 assisted photochemical oxidation of Remazol Turquoise Blue. Chem. Eng. J. 200–202, 373-379. Kalajzic, T., Bianchi, M., Muntau, H., Kettrup, A., 1998. Polychlorinated biphenyls

207

(PCBs) and organochlorine pesticides (OCPs) in the sediments of an Italian drinking water reservoir. Chemosphere 36, 1615-1625. Kang, J.-W., Hoffmann, M.R., 1998. Kinetics and Mechanism of the Sonolytic Destruction of Methyl tert-Butyl Ether by Ultrasonic Irradiation in the Presence of Ozone. Environ. Sci. Technol.32, 3194-3199. Kang, S.-F., Liao, C.-H., Po, S.-T., 2000. Decolorization of textile wastewater by photo- Fenton oxidation technology. Chemosphere 41, 1287-1294. Karagedov, G., Lyakhov, N., 2003. Mechanochemical grinding of inorganic oxides. KONA Powder Part. J. 21, 76-87. Kim, T.-H., Park, C., Shin, E.-B., Kim, S., 2002. Decolorization of disperse and reactive dyes by continuous electrocoagulation process. Desalination 150, 165-175. Kim, Y.-H., Carraway, E.R., 2000. Dechlorination of Pentachlorophenol by Zero Valent Iron and modified Zero Valent Irons. Environ. Sci. Technol.34, 2014-2017. Khan, J.A., He, X., Khan, H.M., Shah, N.S., Dionysiou, D.D., 2013. Oxidative degradation of atrazine in aqueous solution by UV/H2O2/Fe2+, UV//Fe2+ and UV//Fe2+ processes: A comparative study. Chem. Eng. J. 218, 376-383. Khan, J.A., He, X., Shah, N.S., Khan, H.M., Hapeshi, E., Fatta-Kassinos, D.,

Dionysiou, D.D., 2014. Kinetic and mechanism investigation on the

photochemical degradation of atrazine with activated H2O2, S2O82− and

HSO5−. Chemical Engineering Journal 252, 393-403.

Khan, S., He, X., Khan, H.M., Boccelli, D., Dionysiou, D.D., 2016. Efficient

degradation of lindane in aqueous solution by iron (II) and/or UV activated

peroxymonosulfate. Journal of Photochemistry and Photobiology A: Chemistry

316, 37-43.

Kitano, M., Matsuoka, M., Ueshima, M., Anpo, M., 2007. Recent developments in titanium oxide-based photocatalysts. Appl. Catal. A: Gen.325, 1-14. Kochany, J., Bolton, J.R., 1992. Mechanism of photodegradation of aqueous organic

208

pollutants. 2. Measurement of the primary rate constants for reaction of hydroxyl radicals with benzene and some halobenzenes using an EPR spin-trapping method following the photolysis of hydrogen peroxide. Environ. Sci. Technol.26, 262- 265. Kouras, A., Zouboulis, A., Samara, C., Kouimtzis, T., 1998. Removal of pesticides from aqueous solutions by combined physicochemical processes—the behaviour of lindane. Environ. Pollut. 103, 193-202. Kundu, S., Pal, A., Dikshit, A.K., 2005. UV induced degradation of herbicide 2,4-D: kinetics, mechanism and effect of various conditions on the degradation. Sep. Purif. Technol 44, 121-129. Lagunas-Allué, L., Martínez-Soria, M.-T., Sanz-Asensio, J., Salvador, A., Ferronato, C., Chovelon, J.M., 2012. Degradation intermediates and reaction pathway of pyraclostrobin with TiO2 photocatalysis. Appl. Catal. B: Environ.115–116, 285- 293. Lee, B., Jeong, S.W., 2009. Effects of additives on 2,4,6-trinitrotoluene (TNT) removal and its mineralization in aqueous solution by gamma irradiation. J. Hazard. Mater. 165, 435-440. Lee, B., Lee, M., 2005. Decomposition of 2,4,6-trinitrotoluene (TNT) by gamma irradiation. Environ Sci Technol 39, 9278-9285. Legrini, O., Oliveros, E., Braun, A.M., 1993. Photochemical processes for water treatment. Chem. Rev. 93, 671-698. Lewins, J., Becker, M., Kurucz, C., Waite, T., Cooper, W., Nickelsen, M., 1991. High Energy Electron Beam Irradiation of Water, Wastewater and Sludge. Adv. Nucl. Sci. Technol. Springer US, pp. 1-43. Li, S., Elliott, D.W., Spear, S.T., Ma, L., Zhang, W.-X., 2011. Hexachlorocyclohexanes in the environment: Mechanisms of dechlorination. Crit. Rev. Environ. Sci. Technol. 41, 1747-1792. Li, X.-q., Elliott, D.W., Zhang, W.-x., 2006. Zero-Valent Iron Nanoparticles for abatement of environmental pollutants: Materials and Engineering Aspects. Crit. Rev. Solid State Mater. Sci. 31, 111-122.

209

Li, X., Ma, J., Liu, G., Fang, J., Yue, S., Guan, Y., Chen, L., Liu, X., 2012. Efficient reductive dechlorination of Monochloroacetic Acid by Sulfite/UV Process. Environ. Sci. Technol.46, 7342-7349. Li, Y.-F., McMillan, A., Scholtz, M.T., 1996. Global HCH Usage with 1° × 1° Longitude/Latitude Resolution. Environ. Sci. Technol.30, 3525-3533. Li, Y.F., 1999. Global technical hexachlorocyclohexane usage and its contamination consequences in the environment: from 1948 to 1997. Sci. Total Environ. 232, 121-158. Liang, C.J., Bruell, C.J., Marley, M.C., Sperry, K.L., 2003. Thermally activated persulfate oxidation of trichloroethylene (TCE) and 1,1,1- trichloroethane (TCA) in aqueous systems and soil slurries. Soil Sediment. Contam. 12, 207–228. Liang, C., Bruell, C.J., Marley, M.C., Sperry, K.L., 2004. Persulfate oxidation for in situ remediation of TCE. I. Activated by ferrous ion with and without a persulfate- thiosulfate redox couple. Chemosphere 55, 1213-1223. Liang, C., Huang, C.F., Mohanty, N., Kurakalva, R.M., 2008. A rapid spectrophotometric determination of persulfate anion in ISCO. Chemosphere 73, 1540-1543. Li, Y.F., Bidleman, T.F., Barrie, L.A., McConnell, L.L., 1998. Global hexachlorocyclohexane use trends and their impact on the Arctic atmospheric environment. Geophys. Res. Lett. 25, 39-41. Lin, K., Cooper, W.J., Nickelsen, M.G., Kurucz, C.N., Waite, T.D., 1995. Decomposition of aqueous solutions of phenol using high energy electron beam irradiation—A large scale study. Appl. Radiat. Isot. 46, 1307-1316. Lin, S., Su, G., Zheng, M., Jia, M., Qi, C., Li, W., 2011. The degradation of 1,2,4-

trichlorobenzene using synthesized Co3O4 and the hypothesized mechanism.

Journal of hazardous materials 192, 1697-1704.

Liu, S.-Y., Chen, Y.-P., Yu, H.-Q., Zhang, S.-J., 2005. Kinetics and mechanisms of radiation-induced degradation of acetochlor. Chemosphere 59, 13-19. Liu, X., Peng, P.a., Fu, J., Huang, W., 2003. Effects of FeS on the transformation kinetics of γ-Hexachlorocyclohexane. Environ. Sci. Technol.37, 1822-1828.

210

Liu, C.S., Shih, K., Sun, C.X., Wang, F., 2012. Oxidative degradation of propachlor by ferrous and copper ion activated persulfate. Sci Total Environ 416, 507-512. Loiselle, S., Branca, M., Mulas, G., Cocco, G., 1996. Selective mechanochemical dehalogenation of Chlorobenzenes over calcium hydride. Environ. Sci. Technol.31, 261-265. Lopez-Blanco, M.C., Reboreda-Rodriguez, B., Cancho-Grande, B., Simal-Gandara, J., 2002. Optimization of solid-phase extraction and solid-phase microextraction for the determination of alpha- and beta-endosulfan in water by gas chromatography- electron-capture detection. J. Chromatogr. A 976, 293-299. Madhavan, J., Muthuraaman, B., Murugesan, S., Anandan, S., Maruthamuthu, P., 2006. Peroxomonosulphate, an efficient oxidant for the photocatalysed degradation of a textile dye, acid red 88. Sol. Energy Mater. Sol. Cells 90, 1875-1887. Mak, F.T., Zele, S.R., Cooper, W.J., Kurucz, C.N., Waite, T.D., Nickelsen, M.G., 1997. Kinetic modeling of carbon tetrachloride, chloroform and methylene chloride removal from aqueous solution using the electron beam process. Water Res. 31, 219-228. Malato, S., Blanco, J., Vidal, A., Richter, C., 2002. Photocatalysis with solar energy at a pilot-plant scale: an overview. Appl. Catal. B: Environ.37, 1-15. Malato, S., Caceres, J., Agüera, A., Mezcua, M., Hernando, D., Vial, J., Fernández-Alba,

A.R., 2001. Degradation of Imidacloprid in Water by Photo-Fenton and TiO2 Photocatalysis at a Solar Pilot Plant: A Comparative Study. Environ. Sci. Technol. 35, 4359-4366. Malato, S., Blanco, J., Richter, C., Braun, B., Maldonado, M.I., 1998. Enhancement of the rate of solar photocatalytic mineralization of organic pollutants by inorganic oxidizing species. Appl. Catal. B: Environ.17, 347-356. Martinez-Huitle, C.A., Ferro, S., 2006. Electrochemical oxidation of organic pollutants for the wastewater treatment: direct and indirect processes. Chem. Soc. Rev. 35, 1324-1340. Masomboon, N., Ratanatamskul, C., Lu, M.C., 2009. Chemical oxidation of 2,6- dimethylaniline in the fenton process. Environ. Sci. Technol. 43, 8629-8634.

211

Matsumoto, Y., Murakami, M., Shono, T., Hasegawa, T., Fukumura, T., Kawasaki, M., Ahmet, P., Chikyow, T., Koshihara, S.-y., Koinuma, H., 2001. Room-temperature ferromagnetism in transparent transition metal-doped titanium dioxide. Science 291, 854-856. Matta, R., Hanna, K., Chiron, S., 2007. Fenton-like oxidation of 2,4,6-trinitrotoluene using different iron minerals. Sci. Total Environ.385, 242-251. Maurino, V., Calza, P., Minero, C., Pelizzetti, E., Vincenti, M., 1997. Light-assisted 1,4- dioxane degradation. Chemosphere 35, 2675-2688. Mekprasart, W., Pecharapa, W., 2011. Synthesis and characterization of Nitrogen-doped

TiO2 and its photocatalytic activity enhancement under visible light. Energy Procedia 9, 509-514. Metcalf, R.L., 1955. Organic Insecticides: Their Chemistry and Mode of Action. Interscience Publishers. Middeldorp, P.J.M., Jaspers, M., Zehnder, A.J.B., Schraa, G., 1996. Biotransformation of α-, β-, γ-, and δ-Hexachlorocyclohexane under Methanogenic conditions. Environ. Sci. Technol.30, 2345-2349. Mincher, B.J., Brey, R.R., Rodriguez, R.G., Pristupa, S., Ruhter, A., 2002. Increasing PCB radiolysis rates in transformer oil. Radiat. Phys. Chem. 65, 461-465. Mincher, B.J., Meikrantz, D.H., Murphy, R.J., Gresham, G.L., Connolly, M.J., 1991. Gamma-ray induced degradation of PCBs and pesticides using spent reactor fuel. Int. J. Rad. Appl. Instrum: A. Appl. Radiat. Isot. 42, 1061-1066. Mingos, D.M.P., Baghurst, D.R., 1991. Tilden Lecture. Applications of microwave dielectric heating effects to synthetic problems in chemistry. Chem. Soc. Rev.20, 1-47. Mohamed, K.A., Basfar, A.A., Al-Kahtani, H.A., Al-Hamad, K.S., 2009. Radiolytic degradation of malathion and lindane in aqueous solutions. Radiat. Phys. Chem. 78, 994-1000. Moore, N.W., Walker, C.H., 1964. Organic chlorine insecticide residues in wild birds. Nature 201, 1072-1073. Mucka, V., Polakova, D., Pospisil, M., Silber, R., 2003. Dechlorination of chloroform in

212

aqueous solutions influenced by nitrate ions and hydrocarbonate ions. Radiat. Phys. Chem. 68, 787-791. Murov, S.L., Carmichael, I., Hug, G.L., 1993. Handbook of Photochemistry, Second ed. Marcel Decker, New York, pp. 330-336

Nakata, K., Fujishima, A., 2012. TiO2 photocatalysis: Design and applications. J. Photoch. Photobiol. C: Photoch. Rev. 13, 169-189 Nasir, K., Bilto, Y.Y., Al-Shuraiki, Y., 1998. Residues of chlorinated hydrocarbon insecticides in human milk of Jordanian women. Environ Pollut 99, 141-148. National, E.P.A. Priorities List, 2005. Potential for Human Exposure. Frequency of NPL Sites with α -Hexachlorocyclohexane. 183-220.

•− Neta, P., Zemel, I.H., 1977. Rate constants and mechanism of reaction of SO4 with aromatic compounds. J. Am. Chem. Soc., 99, 163-164. Neta, P., Huie, R.E., Ross, A.B., 1988. Rate Constants for Reactions of Inorganic

Radicals in Aqueous Solution. Journal of Physical and Chemical Reference Data

17, 1027-1284.

Neyens, E., Baeyens, J., 2003. A review of classic Fenton’s peroxidation as an advanced oxidation technique. J. Hazard. Mater.98, 33-50. Nfodzo, P., Choi, H., 2011. Triclosan decomposition by sulfate radicals: Effects of oxidant and metal doses. Chem. Eng. J. 174, 629-634. Nickelsen, M.G., Cooper, W.J., Lin, K., Kurucz, C.N., Waite, T.D., 1994. High energy electron beam generation of oxidants for the treatment of benzene and toluene in the presence of radical scavengers. Water Res. 28, 1227-1237. Nienow, A.M., Bezares-Cruz, J.C., Poyer, I.C., Hua, I., Jafvert, C.T., 2008. Hydrogen peroxide-assisted UV photodegradation of Lindane. Chemosphere 72, 1700-1705. Nitoi, I., Oncescu, T., Oancea, P., 2013. Mechanism and kinetic study for the degradation of lindane by photo-Fenton process. J. Ind. Eng. Chem.19, 305-309. Nomura, Y., Fujiwara, K., Terada, A., Nakai, S., Hosomi, M., 2012. Mechanochemical degradation of γ-hexachlorocyclohexane by a planetary ball mill in the presence of CaO. Chemosphere 86, 228-234.

213

Ocampo-Perez, R., Rivera-Utrilla, J., Sanchez-Polo, M., Lopez-Penalver, J.J., Leyva- Ramos, R., 2011. Degradation of antineoplastic cytarabine in aqueous solution by gamma radiation. Chem. Eng. J. 174, 1-8. Oesch, F., Friedberg, T., Herbst, M., Paul, W., Wilhelm, N., Bentley, P., 1982. Effects of lindane treatment on drug metabolizing enzymes and liver weight of CF1 mice in which it evoked hepatomas and in non-susceptible rodents. Chem. Biol. Interact.40, 1-14. Oh, S.Y., Kim, H.W., Park, J.M., Park, H.S., Yoon, C., 2009. Oxidation of polyvinyl alcohol by persulfate activated with heat, Fe2+, and zero-valent iron. J. Hazard. Mater.168, 346-351. Okeke, B.C., Siddique, T., Arbestain, M.C., Frankenberger, W.T., 2002. Biodegradation of gamma-hexachlorocyclohexane (lindane) and alpha-hexachlorocyclohexane in water and a soil slurry by a Pandoraea species. J. Agric. Food Chem 50, 2548- 2555. Ollis, D.F., Al-Ekabi, H., 1992. Photocatalytic purification and treatment of water and

air: proceedings of the 1st International Conference on TiO2 Photocatalytic Purification and Treatment of Water and Air, London, Ontario, Canada, 8-13. Orloff, H.D., 1954. The Stereoisomerism of Cyclohexane Derivatives. Chem. Rev. 54, 347-447. Pagano, M., Volpe, A., Mascolo, G., Lopez, A., Locaputo, V., Ciannarella, R., 2011. Peroxymonosulfate-Co(II) oxidation system for the removal of the non-ionic surfactant Brij 35 from aqueous solution, Chemosphere, 86, 329-334. Pages, N., Sauviat, M.P., Bouvet, S., Goudey-Perriere, F., 2002. Reproductive toxicity of lindane. Journal de la Societe de biologie 196, 325-338. Panagiotou, T., Levendis, Y.A., Carlson, J., Dunayevskiy, Y.M., Vouros, P., 1996. Aromatic hydrocarbon emissions from burning poly(styrene), poly(ethylene) and PVC particles at high temperatures. Combust. Sci. Tech. 116–117, 91– 128. Parra, S., Elena Stanca, S., Guasaquillo, I., Ravindranathan Thampi, K., 2004.

Photocatalytic degradation of atrazine using suspended and supported TiO2.

Applied Catalysis B: Environmental 51, 107-116. 214

Patel, U., Suresh, S., 2008. Electrochemical treatment of pentachlorophenol in water and pulp bleaching effluent. Sep. Purif. Technol. 61, 115-122. Patton, G.W., Hinckley, D.A., Walla, M.D., Bidleman, T.F., Hargrave, B.T., 1989. Airborne organochlorines in the Canadian High Arctic. Tellus B 41B, 243-255. Peternel, I., Koprivanac, N., Kusic, H., 2006. UV-based processes for reactive azo dye mineralization. Water Res. 40, 525-532. Peyton, G.R., Glaze, W.H., 1988. Destruction of pollutants in water with ozone in combination with ultraviolet radiation. 3. Photolysis of aqueous ozone. Environ. Sci. Technol.22, 761-767. PHG 1999, Public Health Goal for Lindane in Drinking Water, California Public Health Goal (PHG). Available at http://www.oehha.org/water/phg/pdf/lindan_f.pdf (downloaded on June 29, 2014) Pignatello, J.J., Oliveros, E., MacKay, A., 2006. Advanced oxidation processes for organic contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev. Environ. Sci. Technol. 36, 1-84. Pignatello, J.J., 1992. Dark and photoassisted iron(3+)-catalyzed degradation of chlorophenoxy herbicides by hydrogen peroxide. Environ. Sci. Technol.26, 944- 951. Pignatello, J.J., Liu, D., Huston, P., 1999. Evidence for an Additional oxidant in the photoassisted Fenton reaction. Environ. Sci. Technol.33, 1832-1839. Pirbazari, M., Badriyha, B., Miltner, R., 1991. GAC Adsorber design for removal of chlorinated pesticides. J. Environ. Eng. 117, 80-100. Pizzigallo, M.D.R., Napola, A., Spagnuolo, M., Ruggiero, P., 2004. Influence of inorganic soil components and humic substances on the mechanochemical removal of pentachlorophenol. J. Mater. Sci. 39, 5455-5459. Plakas, K.V., Karabelas, A.J., 2012. Removal of pesticides from water by NF and RO membranes — A review. Desalination 287, 255-265. Planas, C., Caixach, J., Santos, F.J., Rivera, J., 1997. Occurrence of pesticides in Spanish surface waters. Analysis by high resolution gas chromatography coupled to mass spectrometry. Chemosphere 34, 2393-2406.

215

Poulios, I., Micropoulou, E., Panou, R., Kostopoulou, E., 2003. Photooxidation of eosin Y in the presence of semiconducting oxides. Appl. Catal. B: Environ.41, 345-355. Quintero, J.C., Moreira, M.T., Feijoo, G., Lema, J.M., 2005. Anaerobic degradation of hexachlorocyclohexane isomers in liquid and soil slurry systems. Chemosphere 61, 528-536. Quiroz, M.A., Bandala, E.R., Martínez-huitle, C.A., 2011. Advanced oxidation processes (AOPs) for removal of pesticides from aqueous media. Pesticides - Formulations, Effects, Fate, Prof. Margarita Stoytcheva (Ed.), ISBN: 978-953-307-532-7, Available from: http://www.intechopen.com/books/pesticides-formulations-effects- fate/advanced-oxidation-processes-aops-forremoval-of-pesticides-from-aqueous-media (downloaded on March 13, 2014). Radjenović, J., Farré, M.J., Mu, Y., Gernjak, W., Keller, J., 2012. Reductive electrochemical remediation of emerging and regulated disinfection byproducts. Water Res. 46, 1705-1714. Rahn, R.O., 1997. Potassium Iodide as a Chemical Actinometer for 254 nm Radiation: Use of lodate as an Electron Scavenger. Photochem. Photobiol. 66, 450-455. Rahn, R.O., Stefan, M.I., Bolton, J.R., Goren, E., Shaw, P.S., Lykke, K.R., 2003. Quantum yield of the iodide-iodate chemical actinometer: dependence on wavelength and concentrations. Photochem Photobiol 78, 146-152. Rankenberger, W.I.T.F., 2002. Biodegradation of γ-Hexachlorocyclohexane ( Lindane ) and α–Hexachlorocyclohexane in Water and Soil Slurry by a Pandoraea Species. J. Agric. Food Chem. 50, 2548-2555. Rauf, M.A., Ashraf, S.S., 2009. Radiation induced degradation of dyes--an overview. J. Hazard. Mater. 166, 6-16. Rastogi, A., Al-Abed, S.R., Dionysiou, D.D., 2009. Sulfate radical-based ferrous– peroxymonosulfate oxidative system for PCBs degradation in aqueous and sediment systems. Appl. Catal. B: Environ.85, 171-179. Reddy, V.R., Behera, B., 2006. Impact of water pollution on rural communities: An economic analysis. Ecol. Economics 58, 520-537. Riga, A., Soutsas, K., Ntampegliotis, K., Karayannis, V., Papapolymerou, G., 2007.

216

Effect of system parameters and of inorganic salts on the decolorization and

degradation of Procion H-exl dyes. Comparison of H2O2/UV, Fenton, UV/Fenton,

TiO2/UV and TiO2/UV/H2O2 processes. Desalination 211, 72-86. Rodriguez, S., Santos, A., Romero, A., 2011. Effectiveness of AOP's on abatement of emerging pollutants and their oxidation intermediates: Nicotine removal with Fenton's Reagent. Desalination 280, 108-113. Reints, W., Pratt, D.A., Korth, H.-G., Mulder, P., 2000. O−O Bond dissociation enthalpy

in Di(trifluoromethyl) peroxide (CF3OOCF3) as determined by very low pressure pyrolysis. Density Functional Theory Computations on O−O and O−H Bonds in (Fluorinated) Derivatives. J. Phys. Chem: A 104, 10713-10720. Rivas, F.J., Beltrán, F.J., Carvalho, F., Alvarez, P.M., 2005. Oxone-promoted wet air oxidation of landfill leachates. Ind. Eng. Chem. Res. 44, 749-758. Romero, A., Santos, A., Vicente, F., Gonzalez, C., 2010. Diuron abatement using activated persulphate: Effect of pH, Fe(II) and oxidant dosage. Chem. Eng. J. 162, 257-265. Sacco, O., Vaiano, V., Han, C., Sannino, D., Dionysiou, D.D., 2015.

Photocatalytic removal of atrazine using N-doped TiO2 supported on phosphors.

Applied Catalysis B: Environmental 164, 462-474.

Sadr Ghayeni, S.B., Madaeni, S.S., Fane, A.G., Schneider, R.P., 1996. Aspects of microfiltration and reverse osmosis in municipal wastewater reuse. Desalination 106, 25-29. Sahu, S.K., Patnaik, K.K., Sharmila, M., Sethunathan, N., 1990. Degradation of Alpha-, Beta-, and Gamma-Hexachlorocyclohexane by a soil bacterium under aerobic conditions. Appl. Environ. Microbiol. 56, 3620-3622. Saien, J., Ojaghloo, Z., Soleymani, A.R., Rasoulifard, M.H., 2011. Homogeneous and heterogeneous AOPs for rapid degradation of Triton X-100 in aqueous media via UV light, nano titania hydrogen peroxide and potassium persulfate. Chem. Eng. J. 167, 172-182. Salvador, R., Casal, B., Yates, M., Martıń -Luengo, M.A., Ruiz-Hitzky, E., 2002.

217

Microwave decomposition of a chlorinated pesticide (Lindane) supported on modified sepiolites. Appl. Clay Sci. 22, 103-113. Sang, S., Petrovic, S., Cuddeford V; 1999. Lindane – A Review of Toxicity and Environmental Fate, A Technical report of World Wildlife Fund Canada, 1-65. Internet Site: www.wwf.ca. Sanromán, M.A., Pazos, M., Ricart, M.T., Cameselle, C., 2004. Electrochemical decolourisation of structurally different dyes. Chemosphere 57, 233-239. Sarkar, A., Nagarajan, R., Chaphadkar, S., Pal, S., Singbal, S.Y.S., 1997. Contamination of organochlorine pesticides in sediments from the Arabian Sea along the west coast of India. Water Res. 31, 195-200. Sarkar, S.K., Bhattacharya, B.D., Bhattacharya, A., Chatterjee, M., Alam, A., Satpathy, K.K., Jonathan, M.P., 2008. Occurrence, distribution and possible sources of organochlorine pesticide residues in tropical coastal environment of India: An overview. Environ. Int. 34, 1062-1071. Schlimn, C., Heitz, E., 1996. Development of a wastewater treatment process: Reductive dehalogenation of chlorinated hydrocarbons by metals. Environ. Progr. 15, 38-47. Schmidt, K.H., Han, P., Bartels, D.M., 1995. Radiolytic yields of the hydrated electron from transient conductivity. Improved calculation of the hydrated electron

− diffusion coefficient and analysis of some diffusion-limited eaq reaction rates. J. Phys. Chem. 99, 10530-10539. Schmelling, D., Poster, D., Chaychian, M., Neta, P., McLaughlin, W., Silverman, J., Al- Sheikhly, M., 1998. Applications of ionizing radiation to the remediation of materials contaminated with heavy metals and polychlorinated biphenyls. Radiat. Phys. Chem. 52, 371-377. Schüth, C., Reinhard, M., 1998. Hydrodechlorination and hydrogenation of aromatic compounds over palladium on alumina in hydrogen-saturated water. Appl. Catal. B: Environ.18, 215-221. Senthilnathan, J., Philip, L., 2009. Removal of mixed pesticides from drinking water

system by photodegradation using suspended and immobilized TiO2. J. Environ. Sci. Heal. B 44, 262-270.

218

Senthilnathan, J., Philip, L., 2010. Photocatalytic degradation of lindane under UV and

visible light using N-doped TiO2. Chem. Eng. J. 161, 83-92. Sheoran, M., 2008. Advanced oxidation processes for the degradation of Pesticides. MS Thesis. Department of Biotechnology and Environmental Sciences, Thapar University, Patiala, Punjab, India. Shukla, P., Fatimah, I., Wang, S., Ang, H.M., Tadé, M.O., 2010. Photocatalytic generation of sulphate and hydroxyl radicals using zinc oxide under low-power UV to oxidise phenolic contaminants in wastewater. Catal. Today 157, 410-414. Singh, A., Kremers, W., 2002. Radiolytic dechlorination of polychlorinated biphenyls using alkaline 2-propanol solutions. Radiat. Phys. Chem. 65, 467-472. Singh, R., Misra, V., Singh, R.P., 2011. Remediation of γ-hexachlorocyclohexane contaminated soil using nanoscale zero-valent iron. J. Bionanosci. 5, 82-87. Snyder, S.A., Adham, S., Redding, A.M., Cannon, F.S., DeCarolis, J., Oppenheimer, J., Wert, E.C., Yoon, Y., 2007. Role of membranes and activated carbon in the removal of endocrine disruptors and pharmaceuticals. Desalination 202, 156-181. Somasundaram, L., Coats, J.R., 1991. Pesticide transformation products. Fate

and significance in the environment. ACS Symposium Series.

Song, W., Yan, S., Cooper, W.J., Dionysiou, D.D., O’Shea, K.E., 2012. Hydroxyl radical oxidation of Cylindrospermopsin (Cyanobacterial Toxin) and its role in the photochemical transformation. Environ. Sci. Technol.46, 12608-12615. Sourirajan, S., 1970. Reverse osmosis. Academic Press. Spinks J. W. T, Woods, R.J., 1990. An Introduction to Radiation Chemistry, 3rd edition, John Wiley and Sons Inc., New York.

Stafford, U., Gray, K.A., Kamat, P.V., 1994. Radiolytic and TiO2-assisted photocatalytic degradation of 4-Chlorophenol. A comparative study. J. Phys. Chem. 98, 6343- 6351. Sweeney, E.A., Chipman, J.K., Forsythe, S.J., 1994. Evidence for direct-acting oxidative genotoxicity by reduction products of azo dyes. Environ Health Perspect 102 Suppl 6, 119-122. Taghipour, F., Evans, G.J., 1997. Radiolytic dechlorination of chlorinated organics. 219

Radiat. Phys. Chem. 49, 257-264. Tahri, L., Elgarrouj, D., Zantar, S., Mouhib, M., Azmani, A., Sayah, F., 2010. Wastewater treatment using gamma irradiation: Tétouan pilot station, Morocco. Radiat. Phys. Chem. 79, 424-428. Tamimi, M., Qourzal, S., Barka, N., Assabbane, A., Ait-Ichou, Y., 2008. Methomyl degradation in aqueous solutions by Fenton's reagent and the photo-Fenton system. Sep. Purif. Technol 61, 103-108. Tanabe, S., Iwata, H., Tatsukawa, R., 1994. Global Contamination by Persistent Organochlorines and Their Ecotoxicological Impact on Marine Mammals. Sci. Total Environ.154, 163-177. Tanaka, Y., Zhang, Q., Saito, F., 2003. Mechanochemical dechlorination of Trichlorobenzene on oxide surfaces. J. Phys. Chem. B 107, 11091-11097. Tang, W.Z., Tassos, S., 1997. Oxidation kinetics and mechanisms of trihalomethanes by Fenton's reagent. Water Res. 31, 1117-1125. Thacker, N.P., Vaidya, M.V., Sipani, M., Kalra, A., 1997. Removal technology for pesticide contaminants in potable water. J. Environ. Sci. Heal. B 32, 483-496. Tieyu, W., Yonglong, L., Hong, Z., Yajuan, S., 2005. Contamination of persistent organic pollutants (POPs) and relevant management in China. Environ. Int. 31, 813-821. Traina, M.E., Rescia, M., Urbani, E., Mantovani, A., Macrı̀, C., Ricciardi, C., Stazi, A.V., Fazzi, P., Cordelli, E., Eleuteri, P., Leter, G., Spanò, M., 2003. Long-lasting effects of lindane on mouse spermatogenesis induced by in utero exposure. Reproduct. Toxicol. 17, 25-35. Triantis, T.M., Fotiou, T., Kaloudis, T., Kontos, A.G., Falaras, P., Dionysiou, D.D.,

Pelaez, M., Hiskia, A., 2012. Photocatalytic degradation and mineralization of

microcystin-LR under UV-A, solar and visible light using nanostructured nitrogen

doped TiO2. Journal of hazardous materials 211–212, 196-202.

Tusell, J.M., Suñol, C., Gelpí, E., Rodriguez-Farré, E., 1988. Effect of lindane at repeated low doses. Toxicol. 49, 375-379.

220

Uddin, M.H., Hayashi, S., 2009. Effects of dissolved gases and pH on sonolysis of 2,4- dichlorophenol. J. Hazard. Mater.170, 1273-1276. Ullah, R., Malik, R.N., Qadir, A., 2009. Assessment of groundwater contamination in an industrial city , Sialkot , Pakistan. Afr. J. Environ. Sci. Technol. 3, 429-446. Umebayashi, T., Yamaki, T., Itoh, H., Asai, K., 2002. Band gap narrowing of

titanium dioxide by sulfur doping. Applied Physics Letters 81, 454-456.UNEP (United Nations Environment Programme), 2001. Stockholm convention on persistent organic pollutants, as amended 2009. Geneva, Switzerland. U.S D.H.H.S., 2005, Toxicological profile for alpha-, beta, gamma- and delta- hexachlorocyclohexane, Internet Site: www.atsdr.cdc.gov/toxprofiles/tp43.pdf (downloaded March 10, 2013) Van Doesburg, W., van Eekert, M.H., Middeldorp, P.J., Balk, M., Schraa, G., Stams, A.J., 2005. Reductive dechlorination of beta-hexachlorocyclohexane (beta-HCH) by a Dehalobacter species in coculture with a Sedimentibacter sp. FEMS microbial. ecol. 54, 87-95. Vassilakis, C., Pantidou, A., Psillakis, E., Kalogerakis, N., Mantzavinos, D., 2004. Sonolysis of natural phenolic compounds in aqueous solutions: degradation pathways and biodegradability. Water Res. 38, 3110-3118. Vijgen, J., Abhilash, P.C., Li, Y.F., Lal, R., Forter, M., Torres, J., Singh, N., Yunus, M., Tian, C., Schaffer, A., Weber, R., 2011. Hexachlorocyclohexane (HCH) as new Stockholm Convention POPs--a global perspective on the management of Lindane and its waste isomers. Environ. Sci. Pollut. Res. int. 18, 152-162. Voldner, E.C., Li, Y.-F., 1995. Global usage of selected persistent organochlorines. Sci. Total Environ. 160–161, 201-210. Von Gunten, U., 2003. Ozonation of drinking water: Part I. Oxidation kinetics and product formation. Water Res. 37, 1443-1467. Walker, K., Vallero, D.A., Lewis, R.G., 1999. Factors influencing the distribution of lindane and other hexachlorocyclohexanes in the environment. Environ. Sci. Technol. 33, 4373-4378. Walling, C., Camaioni, D.M., 1975. Aromatic hydroxylation by peroxydisulfate.

221

Journal of the American Chemical Society 97, 1603-1604.

Wallington, T.J., Dagaut, P., Liu, R., Kurylo, M.J., 1988. Gas-phase reactions of hydroxyl radicals with the fuel additives methyl tert-butyl ether and tert-butyl alcohol over the temperature range 240-440 K. Environ. Sci. Technol.22, 842- 844. Wang, P., Yang, S., Shan, L., Niu, R., Shao, X., 2011. Involvements of chloride ion in decolorization of Acid Orange 7 by activated peroxydisulfate or peroxymonosulfate oxidation. J. Environ. Sci. 23, 1799-1807.

Wang, P., Yap, P.-S., Lim, T.-T., 2011. C–N–S tridoped TiO2 for photocatalytic degradation of tetracycline under visible-light irradiation. Appl. Catal. A: Gen.399, 252-261. Wang, T., Waite, T.D., Kurucz, C., Cooper, W.J., 1994. Oxidant reduction and biodegradability improvement of paper mill effluent by irradiation. Water Res. 28, 237-241. Wang, Y.R., Chu, W., 2011. Degradation of a xanthene dye by Fe(II)-mediated activation of Oxone process. J. Hazard. Mater.186, 1455-1461. Wang, Y.R., Chu, W., 2013. Photo-assisted degradation of 2,4,5-trichlorophenol

by Electro-Fe(II)/Oxone® process using a sacrificial iron anode: Performance

optimization and reaction mechanism. Chemical Engineering Journal 215-216,

643-650.

Wang, Z., Peng, P., Huang, W., 2009. Dechlorination of gamma-hexachlorocyclohexane by zero-valent metallic iron. J. Hazard. Mater.166, 992-997. Wania, F., 2003. Assessing the potential of persistent organic chemicals for long-range transport and accumulation in polar regions. Environ. Sci. Technol.37, 1344- 1351. Ware, G.W., Whitacre, D.M., 2004. The pesticide book. MeisterPro Information Resources. Wasiewicz, M., Chmielewski, A.G., Getoff, N., 2006. Radiation-induced degradation of aqueous 2,3-dihydroxynaphthalene. Radiat. Phys. Chem. 75, 201-209. 222

Wei, L., Zhu, H., Mao, X., Gan, F., 2011. Electrochemical oxidation process combined with UV photolysis for the mineralization of nitrophenol in saline wastewater. Sep. Purif. Technol. 77, 18-25. Whitmire, M., Bryan, P., Henry, T., Holbrook, J., Lehmann, P., Mollitor, T., Ohorodnik, S., Reed, D., Wietgrefe, H., 2010. Non-clinical dose formulation analysis method validation and sample analysis. The AAPSJ 12, 628-634. WHO, 2004. Lindane in Drinking-water. Background document for development of WHO Guidelines for Drinking-water Quality. Available at; www.who.int/water_sanitation_health/dwq/chemicals/lindane.pdf (downloaded Feb., 20, 2014) Wilcoxon, J.P., Thurston, T.R., 1998. Photocatalysis using semiconductor nanoclusters. MRS Online Proceedings Library 549. Willett, K.L., Ulrich, E.M., Hites, R.A., 1998. Differential Toxicity and Environmental Fates of Hexachlorocyclohexane isomers. Environ. Sci. Technol.32, 2197-2207. Williams, J.A., Cooper, W.J., Mezyk, S.P., Bartels, D.M., 2002. Absolute rate constants for the reaction of the hydrated electron, hydroxyl radical and hydrogen atom with chloroacetones in water. Radiat. Phys. Chem. 65, 327-334. Wintgens, T., Melin, T., Schäfer, A., Khan, S., Muston, M., Bixio, D., Thoeye, C., 2005. The role of membrane processes in municipal wastewater reclamation and reuse. Desalination 178, 1-11. Wiszniowski, J., D. Robert, D., Surmacz-Gorska, J., Miksch, K., Weber, J.-V., 2003.

Photocatalytic mineralization of humic acids with TiO2: effect of pH, sulfate and chloride anions, Int. J. Photoenergy 5, 69–74. Wong, P.K., Yuen, P.Y., 1996. Decolorization and biodegradation of methyl red by Klebsiella pneumoniae RS-13. Water Res.. 30, 1736-1744. Wooley, D., Zimmer L., Dodge, D and Swanson, K., 1985. Effects of lindane-type insecticide in mammals: unsolved problems, Neurotoxicol. 6, 165-192. Wren, J.C., Glowa, G.A., 2000. A simplified kinetic model for the degradation of 2- butanone in aerated aqueous solutions under steady-state gamma-radiolysis. Radiat. Phys. Chem. 58, 341-356.

223

Wu, H., Li, Q., 2012. Application of mechanochemical synthesis of advanced materials. J. Adv. Ceram. 1, 130-137. Wu, J., Zhang, H., Qiu, J., 2012. Degradation of Acid Orange 7 in aqueous solution by a novel electro/Fe2+/peroxydisulfate process. J. Hazard. Mater.215–216, 138-145. Wu, W.Z., Xu, Y., Schramm, K.W., Kettrup, A., 1997. Study of sorption, biodegradation and isomerization of HCH in stimulated sediment/water system. Chemosphere 35, 1887-1894. Xu, B., Gao, N.Y., Cheng, H., Xia, S.J., Rui, M., Zhao, D.D., 2009. Oxidative

degradation of dimethyl phthalate (DMP) by UV/H2O2 process. J. Hazard. Mater.162, 954-959. Xu, S.C., Zhou, H.D., Wei, X.Y., Lu, J., 1989. The Ph-dependence and effects of the oxidative products of some aromatic-compounds in ozonation under UV Irradiation. Ozone-Sci Eng. 11, 281-296. Xu, L., Yuan, R.X., Guo, Y.G., Xiao, D.X., Cao, Y., Wang, Z.H., Liu, J.S., 2013. Sulfate radical-induced degradation of 2,4,6-trichlorophenol: A de novo formation of chlorinated compounds. Chem. Eng. J. 217, 169-173. Yang, S., Wang, P., Yang, X., Shan, L., Zhang, W., Shao, X., Niu, R., 2010. Degradation efficiencies of azo dye Acid Orange 7 by the interaction of heat, UV and anions with common oxidants: persulfate, peroxymonosulfate and hydrogen peroxide. J. Hazard. Mater.179, 552-558. Yang, X., Cao, C., Erickson, L., Hohn, K., Maghirang, R., Klabunde, K., 2009. Photo-

catalytic degradation of Rhodamine B on C-, S-, N-, and Fe-doped TiO2 under visible-light irradiation. Appl. Catal. B: Environ.91, 657-662. Yu, S., Lee, B., Lee, M., Cho, I.H., Chang, S.W., 2008. Decomposition and mineralization of cefaclor by ionizing radiation: kinetics and effects of the radical scavengers. Chemosphere 71, 2106-2112. Zalazar, C.S., Labas, M.D., Brandi, R.J., Cassano, A.E., 2007. Dichloroacetic acid degradation employing hydrogen peroxide and UV radiation. Chemosphere 66, 808-815. Zaleska, A., Hupka, J., Silowiecki, A., Wiergowski, M., Biziuk, M., 1999.

224

Destruction of chlorinated pesticides in TiO2-enhanced photochemical process.

International Journal of Photoenergy 1(1999) 1-6

Zaleska, A., Hupka, J., Wiergowski, M., Biziuk, M., 2000. Photocatalytic degradation of lindane, p,p '-DDT and methoxychlor in an aqueous environment. J Photoch. Photobio: A 135, 213-220. Zepp, R.G., Faust, B.C., Hoigne, J., 1992. Hydroxyl radical formation in aqueous reactions (pH 3-8) of iron(II) with hydrogen peroxide: the photo-Fenton reaction. Environ. Sci. Technol.26, 313-319. Zhang, J., Zheng, Z., Luan, J., Yang, G., Song, W., Zhong, Y., Xie, Z., 2007. Degradation of hexachlorobenzene by electron beam irradiation. J. Hazard. Mater.142, 431-436. Zhang, J.B., Zheng, Z., Yang, G.J., Zhao, Y.F., 2007. Degradation of microcystin by gamma irradiation. Nucl. Instrum. Meth. A 580, 687-689. Zhang, S.J., Yu, H.Q., Zhao, Y., 2005. Kinetic modeling of the radiolytic degradation of Acid Orange 7 in aqueous solutions. Water Res. 39, 839-846. Zhao, C., Pelaez, M., Dionysiou, D.D., Pillai, S.C., Byrne, J.A., O'Shea, K.E.,

2014. UV and visible light activated TiO2 photocatalysis of 6-hydroxymethyl

uracil, a model compound for the potent cyanotoxin cylindrospermopsin.

Catalysis Today 224, 70-76.

Zhao, J., Zhang, Y., Quan, X., Chen, S., 2010. Enhanced oxidation of 4-chlorophenol using sulfate radicals generated from zero-valent iron and peroxydisulfate at ambient temperature. Sep. Purif. Technol. 71, 302-307. Zhao, X.-K., Yang, G.-P., Wang, Y.-J., Gao, X.-C., 2004. Photochemical degradation of dimethyl phthalate by Fenton reagent. J. Photochem. Photobiol. A: Chem. 161, 215-220. Zheng, T.-C., Richardson, D.E., 1995. Homogeneous aqueous oxidation of organic molecules by Oxone and catalysis by a water-soluble manganese porphyrin complex. Tetrahedron Lett. 36, 833-836.

225

Zinovyev, S.S., Shinkova, N.A., Perosa, A., Tundo, P., 2004. Dechlorination of lindane in the multiphase catalytic reduction system with Pd/C, Pt/C and Raney-Ni. Appl. Catal. B: Environ. 47, 27-36. Zona, R., Solar, S., Gehringer, P., 2002. Degradation of 2,4-dichlorophenoxyacetic acid

by ionizing radiation: influence of oxygen concentration. Water Res 36, 1369-

1374.

226