United States
Environmental Protection
Agency
Office of Water
(4607)
EPA815-R-99-014
April 1999
Alternative Disinfectants
and Oxidants
Guidance Manual
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DISCLAIMER
This manual provides accurate technical data and engineering information on disinfectants and
oxidants that are not as widely used as chlorine. The U.S. Environmental Protection Agency
encourages drinking water treatment utilities and drinking water primacy agencies to examine all
aspects of their current disinfection practices to improve the quality of their finished water
without reducing microbial protection.
This document is EPA guidance only. It does not establish or affect legal rights or obligation.
EPA decisions in any particular case will be made applying the laws and regulation on the basis
of specific facts when permits are issued or regulations promulgated.
Mention of trade names or commercial products does not constitute an EPA endorsement or
recommendation for use.
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ACKNOWLEDGMENTS
The Environmental Protection Agency gratefully acknowledges the assistance of the members of
the Microbial and Disinfection Byproducts Federal Advisory Committee and Technical Working
Group for. their comments and suggestions to improve this document. EPA also wishes to thank
the representatives of drinking water utilities, researches, and the American Water Works
Association for their review and comment. Finally, the EPA would like to recognize the
following individuals for their contribution to this guidance manual:
Don Gates, PhD, DCG, Inc.
Stuart Krasner, Metropolitan Water District of Southern California
Rip Rice, PhD, Rice International Consulting Ent.
Mark LeChavallier, PhD, American Water Works Services Co., Inc.
Jack DeMarco, City of Cincinnati
Blake Atkins, USEPA Region VI
Bob Clement, USEPA Region VIII
Maggie Javdan, USEPA, OGWDW
Mike Schock, USEPA, National Risk Management Research Laboratory
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CONTENTS
1. INTRODUCTION 1-1
1.1 OBJECTIVE OF THIS MANUAL 1-1
1.2 BACKGROUND 1-2
1.3 REGULATORY CONTEXT 1-3
1.3.1 Disinfection Profiling and Benchmarking 1-7
1.4 USE OF DISINFECTANTS AS CHEMICAL OXIDANTS 1-8
1.5 How CHLORINE is ADDRESSED IN THIS GUIDANCE MANUAL 1-8
1.6 A SUMMARY OF ALTERNATIVE DISINFECTANT PROPERTIES 1-9
1.7 SELECTING A DISINFECTION STRATEGY 1-11
1.7.1 Disinfection Strategy Evaluation 1-12
1.7.2 Summary 1-17
1.8 REFERENCES 1-18
2. DISINFECTANT USE IN WATER TREATMENT 2-1
2.1 NEED FOR DISINFECTION IN WATER TREATMENT 2-1
2.1.1 Pathogens of Primary Concern 2-2
2.1.2 Recent Waterbome Outbreaks 2-8
2.1.3 Mechanism of Pathogen Inactivation 2-9
2.2 OTHER USES OF DISINFECTANTS IN WATER TREATMENT 2-10
2.2.1 Minimization of DBP Formation 2-10
2.2.2 Control of Nuisance Asiatic Clams and Zebra Mussels 2-11
2.2.3 Oxidation of Iron and Manganese 2-13
2.2.4 Prevention of Regrowth in the Distribution System and Maintenance of
Biological Stability 2-14
2.2.5 Removal of Taste and Odors Through Chemical Oxidation 2-15
2.2.6 Improvement of Coagulation and Filtration Efficiency 2-15
2.2.7 Prevention of Algal Growth in Sedimentation Basins and Filters 2-16
2.2.8 Removal of Color 2-16
2.3 DISINFECTION BYPRODUCTS AND DISINFECTION RESIDUALS 2-16
2.3.1 Types of DBPs and Disinfection Residuals 2-16
2.3.2 Disinfection Byproduct Formation 2-19
2.3.3 DBP Control Strategies 2-23
2.3.4 CT Factor :..: 2-25
2.4 PATHOGEN INACTIVATION VERSUS DBP FORMATION 2-26
2.5 DISINFECTANT RESIDUAL REGULATORY REQUIREMENTS 2-27
2.6 SUMMARY OF CURRENT NATIONAL DISINFECTION PRACTICES 2-28
2.7 CHLORINE 2-30
2.7.1 Chlorine Chemistry 2-31
2.7.2 Chlorine Generation 2-32
2.7.3 Primary Uses and Points of Application of Chlorine 2-33
2.7.4 Pathogen Inactivation and Disinfection Efficacy 2-35
2.7.5 DBP Formation and Control 2-39
2.7.6 Operational Considerations 2-41
2.8 SUMMARY 2-42
2.8.1 Advantages and Disadvantages of Chlorine Use 2-42
2.8.2 Summary Table 2-44
2.8.3 Reference for Additional Information on Chlorine 2-44
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2.9 REFERENCES 2-45
3. OZONE 3-1
3.1 OZONE CHEMISTRY 3-1
3.2 OZONE GENERATION 3-4
3.2.1 Ozone Production 3-4
3.2.2 System Components 3-5
3.2.3 Operation and Maintenance 3-15
3.3 PRIMARY USES AND POINTS OF APPLICATION OF OZONE 3-16
3.3.1 Primary Uses of Ozone 3-16
3.3.2 Points of Application 3-18
3.3.3 Impact on Other Treatment Processes 3-19
3.3.4 Biologically Active Filtration 3-19
3.3.5 Pathogen Inactivation and Disinfection Efficacy 3-21
3.3.6 Inactivation Mechanisms '. 3-22
3.3.7 Disinfection Parameters 3-22
3.3.8 Inactivation of Microorganisms 3-24
3.4 OZONATION DISINFECTION BYPRODUCTS ; 3-27
3.4.1 Ozone Byproduct Control 3-30
3.5 STATUS OF ANALYTICAL METHODS 3-31
3.5.1 Monitoring of Gas Phase Ozone 3-31
3.5.2 Monitoring of Liquid Phase Residual Ozone 3-35
3.5.3 Bromate Monitoring for Systems Using Ozone 3-37
3.6 OPERATIONAL CONSIDERATIONS 3-38
3.6.1 Process Considerations 3-38
3.6.2 Space Requirements 3-38
3.6.3 Material Selection , 3-39
3.6.4 Ozone System Maintenance 3-39
3.6.5 Ozone Safety 3-39
3.7 SUMMARY ; 3-41
3.7.1 Advantages and Disadvantages of Ozone Use 3-41
3.7.2 Summary Table 3-42
3.8 REFERENCES 3-43
4. CHLORINE DIOXIDE 4-1
4.1 CHLORINE DIOXIDE CHEMISTRY 4-1
4.1.1 Oxidation Potential ....4-1
4.2 GENERATION 4-2
4.2.1 Introduction 4-2
4.2.2 Chlorine Dioxide Purity 4-3
4.2.3 Methods of Generating Chlorine Dioxide 4-4
4.2.4 Generator Design 4-9
4.2.5 Chemical Feed Systems 4-11
4.2.6 Generator Power Requirements 4-13
4.3 PRIMARY USES AND POINTS OF APPLICATION FOR CHLORINE DIOXIDE 4-13
4.3.1 Disinfection 4-13
4.3.2 Taste and Odor Control 4-14
4.3.3 Oxidation of Iron and Manganese 4-14
4.4 PATHOGEN INACTIVATION AND DISINFECTION EFFICACY 4-15
4.4.1 Inactivation Mechanisms 4-15
4.4.2 Environmental Effects , 4-15
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4.4.3 Disinfection Efficacy [[[ '..:. ......... . ..... ................ 4-17
4.5 CHLORINE DIOXIDE DISINFECTION BYPRODUCTS ................... . ........................ . ....... ..... 4-22
4.5.1 Production of Chlorite and Chlorate [[[ 4-22
4.5.2 Organic DBFs Produced by Chlorine Dioxide ................................................. 4-25
4.5.3 Chlorine Dioxide DBP Control Strategies ...................... . ................................. 4-25
4.6 STATUS OF ANALYTICAL METHODS ............. . ........................... : .................... . ............... 4-26
4.6.1 Chlorine Dioxide and Chlorite Analytical Methods ....................... . ................. 4-27
4.6.2 Chlorine Dioxide Monitoring for Systems Using Chlorine Dioxide ................ 4-27
4.6.3 Chlorite Monitoring for Systems Using Chlorine Dioxide ........... ..... ............... 4-28
4.7 OPERATIONAL CONSIDERATIONS [[[ . ............... 4-28
4.7.1 Process Considerations [[[ . ............... 4-30
4.7.2 Generator Operation [[[ . ............ ................ 4-31
4.7.3 Feed Chemicals .................. . ................... . ....... .................................................. 4-31
4.8 SUMMARY [[[ . ............... . ..................................... 4-33
4.8.1 Advantages and Disadvantages of Chlorine Dioxide Use ................................ 4-33
4.8.2 Summary Table [[[ 4-34
4.9 REFERENCES .......... . ......................................... ................. . ......... '. ................... ............... 4-35
5. POTASSIUM PERMANGANATE ...................... . .......................................... . ......................... 5-1
5.1 POTASSIUM PERMANGANATE CHEMISTRY .................... . ..................... . ........................... 5-1
5.1.1 Oxidation Potential ................................ . [[[ 5-1
5.1.2 Ability to Form a Residual ... .......................................... : .................................... 5-1
5.2 GENERATION [[[ . ................................ 5-1
5.3 PRIMARY USES AND POINTS OF APPLICATION .................................... . ........................... 5-2
5.3.1 Primary Uses ............................. . [[[ 5-2
5.3.2 Points of Application [[[ . .................................... 5-4
5.4 PATHOGEN INACTIVATION AND DISINFECTION EFFICACY.. ............................................ 5-4
5.4.1 Inactivation Mechanisms .................................. . ................................................. 5-4
5.4.2 Environmental Effects .......................................... . ............................................. 5-5
5.4.3 Use as a Disinfectant [[[ 5-5
5.5 DISINFECTION BYPRODUCT FORMATION [[[ 5-7
5.5.1 Chapel-Hill and Durham, North Carolina Water Treatment Plants ................... 5-7
5.5.2 American Water Works Association Research Foundation TTHM Study ......... 5-9
5.6 STATUS OF ANALYTICAL METHODS ............................. ............................ ................ . ...... 5-9
5.7 OPERATIONAL CONSIDERATIONS ........................................ ............................................ 5-9
5.8 SUMMARY [[[ . .............. . .................................. 5-10
5.8.1 Advantages and Disadvantages of Potassium Permanganate Use .................... 5-10
5.8.2 Summary Table ........ . [[[ 5-1 1
5.9 REFERENCES ................. [[[ . ........................... . ............... 5-12
6. CHLORAMINES ........................... . [[[ ... ................................ 6-1
6.1 CHLORAMINES CHEMISTRY .................................. .............. . ....................... . .................... 6-1
6.1.1 Equilibrium, Kinetic, and Physiochemical Properties ........................................ 6-1
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CONTENTS
6.4.2 Environmental Effects 6-12
6.4.3 Disinfection Efficacy 6-13
6.5 DBF FORMATION 6-15
6.6 STATUS OF ANALYTICAL METHODS 6-16
6.6.1 Monitoring of Chloramines 6-16
6.6.2 Disinfectant Interferences 6-17
6.6.3 Chloramine Monitoring for Systems Using Chloramines 6-19
6.7 OPERATIONAL CONSIDERATIONS 6-20
6.7.1 Conversion to Chloramination from Chlorination 6-20
6.7.2 Potential Operational Impacts from Chloramination Disinfection 6-22
6.7.3 Special Considerations for Chloramination Facilities 6-25
6.8 SUMMARY 6-27
6.8.1 Advantages and Disadvantages of Chloramine Use 6-27
6.8.2 Summary Table 6-28
6.9 REFERENCES 6-29
7. PEROXONE (OZONE/HYDROGEN PEROXIDE) 7-1
7.1 PEROXONE CHEMISTRY 7-1
7.1.1 Oxidation Reactions 7-2
7.1.2 Reactions with Other Water Quality Parameters 7-3
7.1.3 Comparison between Ozone and Peroxone 7-3
7.2 GENERATION 7-4
7.3 PRIMARY USES AND POINTS OF APPLICATION 7-5
7.3.1 Primary Uses 7-5
7.3.2 Points of Application 7-6
7.4 PATHOGEN INACTIVATION : 7-6
7.4.1 Inactivation Mechanism 7-6
7.4.2 Environmental Effects 7-7
7.4.3 Disinfection Efficacy and Pathogen Inactivation 7-8
7.5 DISINFECTION BYPRODUCTS 7-9
7.6 STATUS OF ANALYTICAL METHODS 7-10
7.6.1 Monitoring of Hydrogen Peroxide 7-10
7.7 OPERATIONAL CONSIDERATIONS 7-13
7.7.1 Process Considerations 7-13
7.7.2 Space Requirements 7-13
7.7.3 Materials 7-13
7.8 SUMMARY 7-14
7.8.1 Advantages and Disadvantages of Peroxone Use (Ozone/Hydrogen Peroxide)7-14
7.8.2 Summary Table 7-15
7.9 REFERENCES 7-16
8. ULTRAVIOLET RADIATION „ 8-1
8.1 UV CHEMISTRY (PHOTOCHEMICAL) 8-1
8.1.1 UV Radiation 8-1
8.1.2 UV Disinfection Reactions 8-2
8.1.3 Process Variables 8-2
8.2 GENERATION 8-3
8.2.1 UV Lamps 8-4
8.2.2 Ballasts , 8-4
8.2.3 UV Reactor Design 8-5
8.3 PRIMARY USES AND POINTS OF APPLICATION 8-8
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8.4 PATHOGEN INACTIVATION AND DISINFECTION EFFICIENCY 8-8
8.4.1 Inactivation Mechanism 8-8
8.4.2 Environmental Effects 8-10
8.4.3 Disinfection Efficacy 8-12
8.5 DISINFECTION BYPRODUCTS OF UV RADIATION 8-15
8.5.1 Ground Water 8-15
8.5.2 Surface Water 8-16
8.5.3 DBF Formation with Chlorination and Chloramination following UV Radiation..
8-16
8.6 STATUS OF ANALYTICAL METHODS 8-17
8.6.1 Monitoring of Generated Ultraviolet Radiation 8-17
8.6.2 Disinfectant Interferences 8-18
8.7 OPERATIONAL CONSIDERATIONS 8-18
8.7.1 Equipment Operation 8-19
8.7.2 Equipment Maintenance 8-19
8.7.3 Standby Power 8-20
8.8 SUMMARY TABLE 8-21
8.9 REFERENCES 8-21
9. COMBINED DISINFECTANTS 9-1
9.1 PRIMARY AND SECONDARY DISINFECTANTS 9-1
9.1.1 DBP Formation with Various Primary and Secondary Disinfectant
Combinations 9-3
9.1.2 Impact of Modifying Disinfection Practices 9-6
9.1.3 Chlorine/Chlorine to Chlorine/Chloramine 9-8
9.1.4 Chlorine/Chlorine to Ozone/Chlorine 9-8
9.1.5 Chlorine/Chlorine to Ozone/Chloramine 9-9
9.1.6 Chlorine/Chlorine to Chlorine Dioxide/Chlorine 9-9
9.1.7 Chlorine/Chlorine to Chlorine Dioxide/Chlorine Dioxide 9-10
9.1.8 Chlorine/Chloramine to Ozone/Chloramine 9-10
9.1.9 Chlorine/Chloramine to Chlorine Dioxide/Chloramine 9-10
9.1.10 Ozone/Chlorine to Ozone/Chloramine 9-11
9.1.11 Summary 9-11
9.2 PATHOGEN INACTIVATION WITH INTERACTIVE DISINFECTANTS 9-11
9.2.1 Inactivation Mechanism 9-12
9.2.2 Environmental Effects 9-13
9.2.3 Pathogen Inactivation Efficiency Using Interactive Disinfectants 9-16
9.2.4 Summary: Pathogen Inactivation with Interactive Disinfectants 9-23
9.3 ANALYTICAL METHODS 9-23
9.3.1 Ozone 9-23
9.3.2 Chlorine Dioxide 9-24
9.3.3 Potassium Permanganate 9-24
9.3.4 Chloramine 9-24
9.3.5 Hydrogen Peroxide 9-24
9.3.6 UV Radiation 9-24
9.3.7 Summary of Analytical Methods 9-24
9.4 SUMMARY 9-25
9.5 REFERENCES 9-26
APPENDIX A. SUMMARY OF DISINFECTANT USAGE IN THE UNITED STATES A-l
APPENDIX B. SELECTED COSTS OF ALTERNATIVE DISINFECTION SYSTEMS B-l
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FIGURES
Figure 1-1. Flow Diagram to Evaluate Current Disinfection Practices 1-14
Figure 1-2. Flow Diagram to Narrow Selection of a New Primary Disinfectant 1-16
Figure 1-3. Flow Diagram to Narrow Selection of a New Secondary Disinfectant 1-17
Figure 2-1. Free Chlorine Giardia and Virus CT Requirements 2-38
Figure 2-2. CT Values for Inactivation of Giardia Cysts by Free Chlorine at 10°C
(at C12 dose of 3.0 mg/L) 2-38
Figure 2-3. CT Values for Inactivation of Giardia Cysts by Free Chlorine at pH 7.0
(at C12 dose of 3.0 mg/L) 2-39
Figure 3-1. Oxidation Reactions of Compounds (Substrate) During Ozonation of Water 3-2
Figure 3-2. Reaction of Ozone and Bromide Ion Can Produce Bromate Ion and Brominated
Organics 3-3
Figure 3-3. Basic Ozone Generator 3-4
Figure 3-5. Schematic of an Air Preparation System 3-7
Figure 3-6. Cylindrical Electrode Schematic 3-9
Figure 3-7. Ozone Bubble Contactor 3-11
Figure 3-8. Sidestream Ozone Injection System 3-12
Figure 3-9. Turbine Mixer Ozone Contactor 3-14
Figure 3-10. CT Values for Inactivation of Viruses by Ozone (pH 6 to 9) 3-26
Figure 3-11. Principal Reactions Producing Ozone Byproducts 3-28
Figure 3-12. Main Pathways of Bromate Ion Formation when Ozone Reacts with Bromide
Ion 3-30
Figure 4-1. Conventional Chlorine Dioxide Generator When Using Chlorine-Chlorite
Method 4-7
Figure 4-2. Chlorine Dioxide Generation Using Recycled Aqueous Chlorine Method 4-8
Figure 4-3. Effect of Temperature on N. Gruberi Cyst Inactivation at pH 7 4-17
Figure 4-4. Comparison of Germicidal Efficiency of Chlorine Dioxide and Chlorine 4-19
Figure 4-5. CT Values for Inactivation of Giardia Cysts by Chlorine Dioxide 4-21
Figure 4-6. CT Values for Inactivation of Viruses by Chlorine Dioxide 4-22
Figure 4-7. C. parvum Inactivation by Chlorine Dioxide at 20°C 4-23
Figure 4-8. C. parvum Inactivation by Chlorine Dioxide at 10°C 4-23
Figure 6-1. Theoretical Breakpoint Curve 6-2
Figure 6-2. Distribution Diagram for Chloramine Species with pH 6-3
Figure 6-3. Gaseous Chlorine Feed System 6-5
Figure 6-4. Hypochlorite Feed System 6-6
Figure 6-5. Anhydrous Ammonia Direct Feed System 6-7
Figure 6-6. Anhydrous Ammonia Solution Feed System 6-8
Figure 6-7. Aqua Ammonia Feed System 6-9
Figure 8-1. Closed Vessel Ultraviolet Reactor 8-6
Figure 8-2. Germicidal Inactivation by UV Radiation 8-9
Figure 8-3. UV Dose Required for Inactivation of MS-2 Coliphage 8-11
Figure 8-4. Particle Interactions that Impact UV Effectiveness 8-12
Figure 8-5. UV Doses Required to Achieve Inactivation of Giardia lamblia Cysts Obtained
from Two Different Sources 8-14
Figure 8-6. Impact of Growth Stage of A. rhysodes on the Required UV Dosage to Achieve
Inactivation 8-14
Figure 9-1. Inactivation of C. parvum Attributed to Synergistic Effects. Application of
Ozone Followed by Chlorine Dioxide 9-14
Figure 9-2. Inactivation of C. parvum Attributed to Synergistic Effects. Application of
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Chlorine Dioxide Followed by Free Chlorine 9-14
Figure 9-3. Inactivation of C. parvum Attributed to Synergistic Effects. Application of
Chlorine Dioxide Followed by Monochloramine 9-15
Figure 9-4. Inactivation of E. coli Using Free Chlorine and Monochloramine 9-18
TABLES
Table 1-1. Key Dates for Regulatory Activities 1-4
Table 1-2. Primary Drinking Water Regulations Related to Microbiological
Contaminants 1-6
Table 1-3. Primary Drinking Water Regulations Related to Disinfection Byproducts 1-6
Table 1-4. Primary Drinking Water Regulations Related to Residual Disinfectants 1-7
Table 1-5. Log Removal/Inactivation through Filtration and Disinfection Required
Under the SWTR 1-7
Table 1-6. Summary of Disinfectant Properties (Based on Typical Disinfectant
Application) ;, 1-11
Table 2-1. Waterborne Diseases from Bacteria 2-3
Table 2-2. Waterborne Diseases from Human Enteric Viruses .2-4
Table 2-3. Waterborne Diseases from Parasites 2-6
Table 2-4. Attributes of the Three Waterborne Pathogens of Concern in Water
Treatment 2-7
Table 2-5. Human Parasitic Protozoans ...-. 2-7
Table 2-6. The Effects of Various Oxidants on Mortality of the Asiatic Clam
(Corbicula flumined) 2-12
Table 2-7. Oxidant Doses Required for Oxidation of Iron and Manganese 2-14
Table 2-8. List of Disinfection Byproducts and Disinfection Residuals 2-17
Table 2-9. Status of Health Information for Disinfectants and DBFs 2-18
Table 2-10. Conditions of Formation of DBFs '. 2-22
Table 2-11. Inorganic DBPs Produced During Disinfection 2-23
Table 2-12. Required Removal of TOC by Enhanced Coagulation for Surface Water
Systems Using Conventional Treatment (percent reduction) 2-25
Table 2-13. CT Values for Inactivation of Viruses 2-26
Table 2-14. CT Values for Inactivation of Giardia Cysts 2-26
Table 2-15. Summary of Disinfection Impacts 2-27
Table 2-16. Disinfection Practices of Water Systems that Include Some Form of
Treatment 2-29
Table 2-17. Ozone Application in Water Treatment Plants in the United States 2-30
Table 2-18. Chlorine Uses and Doses ; 2-34
Table 2-19. Typical Chlorine Points of Application and Uses 2-34
Table 2-20. Typical Chlorine Dosages at Water Treatment Plants 2-35
Table 2-21. Percent Reduction in DBP Formation by Moving Chlorination Point Later
In Treatment Train 2-40
Table 2-22. Summary of Chlorine Disinfection 2-44
Table 3-1. Types of Compressors Used in Air Preparation Systems 3-7
Table 3-2. Comparison of Air and High Purity Oxygen Feed Systems 3-8
Table 3-3. Comparison of Primary Characteristics of Low, Medium, and High Frequency
Ozone Generators ., 3-10
Table 3-4. Bubble Diffuser Contactor Advantages and Disadvantages 3-12
Table 3-5. Injection Contacting Advantages and Disadvantages 3-13
Table 3-6. Turbine Mixer Contactor Advantages and Disadvantages 3-14
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Table 3-7. Criteria for Selecting Ozone Feed Points for Small Systems 3-18
Table 3-8. Summary of Reported Ozonation Requirements for 99 Percent Inactivation
of Cryptosporidium Oocycts 3-27
Table 3-9. Principal Known Byproducts of Ozonation 3-28
Table 3-10. Characteristics and Comparisons of Gas-Phase Ozone Analytical Methods 3-33
Table 3-11. Characteristics and Comparisons of Residual Ozone Analytical Methods 3-36
Table 3-12. Summary of Ozone Disinfection Considerations 3-42
Table 4-1. Commercial Chlorine Dioxide Generators 4-5
Table 4-2. Surface Water Chlorine Dioxide Demand Study Results 4-14
Table 4-3. Analytical Methods for Chlorine Dioxide and Related Compounds 4-30
Table 4-4. Properties of Sodium Chlorite as Commercially Available 4-32
Table 4-5. Summary for Chlorine Dioxide 4-34
Table 5-1. Potassium Permanganate CT Values for 2-log Inactivation of MS-2
Bacteriophage 5-7
Table 5-2. Summary of Potassium Permanganate Use 5-12
Table 6-1. Chlorine Dose Required for NH3 - C12 Reaction 6-3
Table 6-2. Time to 99 Percent Conversion of Chlorine to Monochloramine 6-4
Table 6-3. Methods of Chlorine Addition : 6-4
Table 6-4. CT Values for Giardia Cyst Inactivation Using Chloramines 6-15
Table 6-5. CT Values for Virus Inactivation Using Chloramines 6-15
Table 6-6. Characteristics and Comparisons of Monochloraminea Analytical
Methods 6-18
Table 6-7. Survey of Chloramine Users in the United States 6-25
Table 6-8. Summary of Chloramine Disinfection 6-29
Table 7-1. Comparison Between Ozone and Peroxone Oxidation 7-4
Table 7-2. Calculated CT Values (Mg-Min/L) for the Inactivation of Giardia Muris 7-8
Table 7-3. Characteristics and Comparisons of Hydrogen Peroxide Analytical
Methods '. 7-12
Table 7-4. Summary of Peroxone Disinfection Consideration : 7-15
Table 8-1. Water Quality and Associated UV Measurements '..'. 8-3
Table 8-2. Doses Required for MS-2 Inactivation ...„ 8-13
Table 8-3. Summary of UV Disinfection 8-21
Table 9-1. Potential Primary Disinfectants 9-2
Table 9-2. Primary/Secondary Disinfectant Combinations and Typical Applications in
Water Treatment , 9-3
Table 9-3. DBPs Associated with Various Combined Oxidation/Disinfection Processes 9-4
Table 9-4. Strategies for Primary and Secondary Disinfectants 9-6
Table 9-5. Impacts of Disinfection Practice on DBF Formation 9-7
Table 9-6. Virus Inactivation By Individual Disinfectants and Simultaneous
Chloramination 9-17
Table 9-7. C. parvurn Inactivation Using Ozone Followed by Chlorine Dioxide 9-19
Table 9-8. C. parvum Inactivation Using Chlorine Dioxide Followed by Free Chlorine 9-19
Table 9-9. G. muris Inactivation Using Chlorine Dioxide Followed by Free Chlorine 9-19
Table 9-10. B. cereus Inactivation Using Chlorine Dioxide Followed by Free Chlorine 9-19
Table 9-11. C. parvum Inactivation Using Chlorine Dioxide Followed by Chloramine 9-20
Table 9-12. G. muris Inactivation Using Chlorine Dioxide Followed by Chloramine.... 9-20
Table 9-13. G. muris Inactivation Using Ozone Followed by Free Chlorine 9-21
Table 9-14. B. cereus Inactivation Using Chlorine Dioxide Followed by Free Chlorine 9-21
Table 9-15. G. muris Inactivation Using Ozone Followed by Chloramine 9-21
Table 9-16. G. muris Inactivation by Free Chlorine Followed by Monochloramine 9-22
Table 9-17. C. parvum Inactivation by Sequential Application of Ozone and Chloramine 9-22
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Table 9-18. Summary of Combined Disinfectants : 9-25
ACRONYMS
AECL Alternate enhanced coagulant level
ACUK Acid chrome violet K
AOC Assimilable organic carbon
ASDWA Association of State Drinking Water Administrators
AWWA American Water Works Association
AWWARF AWWA Research Foundation
BAG Biologically active carbon
BAF Biologically active filtration
BAT Best Available Technology
BCAA Bromochloroacetic acid
BDOC Biodegradable organic carbon
BMP : Best management practice
BOM Biodegradable Organic Matter (=BDOC + AOC)
Br- Bromide ion
BrO2- Bromite ion
BrO3- Bromate ion
CI Confidence interval
C12 Chlorine
C1O2 Chlorine Dioxide
CT Concentration-Time
CWS Community Water System
D/DBP Disinfectants/disinfection byproducts
DBPR Disinfectants/disinfection byproducts rule
DBP Disinfection byproduct '.'".,,"
DBPFP Disinfection byproduct formation potential
DBPP Disinfection byproduct precursors . ".. „
DBPRAM DBP Regulatory Assessment Model
DBPs Disinfection byproducts ,-.... .
DOC Dissolved organic carbon , > •
DPD N,N-diethyl-p-phenylenediamine , .
DWEL Drinking Water Equivalent Level
EBCT Empty bed contact time
EMSL EPA Environmental Monitoring and Support Laboratory (Cincinnati)
EPA United States Environmental Protection Agency
ESWTR Enhanced Surface Water Treatment Rule
FBR Filter Backwash Rule
FY Fiscal year
GAC Granular activated carbon
GWR Ground Water Rule
GWSS Ground Water Supply Survey
H2O2 Hydrogen Peroxide -
HAAS Haloacetic acids (five)
HOBr Hypobromous acid
HOC1 Hypochlorous acid
1C Ion chromotography
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ICR Information Collection Rule
IESWTR Interim Enhanced Surface Water Treatment Rule
IOA International Ozone Association
IOC Inorganic chemical
KMnO4 Potassium permanganate
LOAEL Lowest observed adverse effect level
LOQ Limit of quantitation
LT1ESWTR Long Term Stage 1 Enhanced Surface Water Treatment Rule
M-DBP Microbial and disinfection byproducts
MCL Maximum Contaminant Level
MCLG Maximum Contaminant Level Goal
MDL Method Detection Limit
mg/L Milligrams per liter
mgd Million gallons per day
MIB Methylisoborneol
MRDL Maximum Residual Disinfectant Level (as mg/1)
MRDLG Maximum Residual Disinfectant Level Goal
MRL Minimum Reporting Level
MX 3-chloro-4-(dichloromethyl)-5-hydroxyl-2(5H)-furanone
NaCl Sodium chloride
NCI National Cancer Institute
ND Not detected
NH2CI Monochloramine
NIPDWR National Interim Primary Drinking Water Regulation
NIOSH National Institute for Occupational Safety and Health
NOAEL No Observed Adverse Effect Level
NOM Natural Organic Matter
NOMS National Organic Monitoring Survey
NORS National Organics Reconnaissance Survey for Halogenated Organics
NPDWR National Primary Drinking Water Regulation
NTNCWS Nontransient noncommunity water system
NTP Normal Temperature and Pressure
O2 Oxygen
O3 Ozone
OBr- Hypobromite ion
OC1- Hypochlorite ion
PCE Perchloroethylene
PE Performance evaluation
POE Point-of-Entry Technologies
POU Point-of-Use Technologies
ppb Parts per billion
ppm Parts per million
PQL Practical Quantitation Level
PTA Packed Tower Aeration
PWS Public water system
RIA Regulatory Impact Analysis
RMCL Recommended Maximum Contaminant Level
RNDB Regulations Negotiation Data Base
RSC Relative Source Contribution
SDWA Safe Drinking Water Act, or the "Act," as amended in 1996
SM Standard Method
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SMCL Secondary Maximum Contaminant Level
SMR Standardized mortality ratios
SOC Synthetic Organic Chemical
SWTR Surface Water Treatment Rule
TCE Trichloroethylene
THM Trihalomethane
THMFP Trihalomethane formation potential
TMV Tobacco mosaic virus
TOC Total organic carbon
TTHM Total trihalomethanes
TWO Technologies Working Group
UV Ultraviolet
VOC Volatile Organic Chemical
WIDE Water Industry Data Base
WS Water supply
XDBPs Halogenated DBFs
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1, INTRODUCTION
1.1 Objective of this Manual
Chlorine is, by far, the most commonly used disinfectant in the drinking water treatment industry
(Sawyer et al., 1994). Today, chlorine is used as a primary disinfectant in the vast majority of all
surface water treatment plants, being used as a pre-disinfectant in more than 63 percent and as a post-
disinfectant in more than 67 percent of all surface water treatment plants (USEPA, 1997). This
manual is organized to provide technical data and engineering information on disinfectants that are
not as widely used as chlorine. Also, where applicable, this document describes the use of these
disinfectants as oxidants and any associated implications.
The U.S. Environmental Protection Agency (EPA) encourages utilities to examine all aspects of their
current disinfection practices to identify opportunities to improve the quality of the finished water
without reducing microbial protection. The objective of this guidance manual is to describe
alternative disinfectants and disinfection techniques that may be used to comply with both the Stage
1 Disinfectants and Disinfection Byproducts Rule (DBPR) and Interim Enhanced Surface Water
Treatment Rule (IESWTR) and highlight advantages and disadvantages of their use.
EPA is not recommending that utilities employ the disinfectants and oxidants discussed in this
manual, nor is it advocating that utilities switch from one disinfectant or oxidant to another. EPA
acknowledges that selection of the most appropriate disinfection technique is a site-specific decision
best left to utility personnel and state agencies. Utilities should use this guidance as an information
resource to assist in the selection of appropriate disinfectants and disinfectant schemes to meet their
specific goals. Extensive bench and/or pilot scale testing and a thorough review of regulatory
requirements should precede changes to disinfection practice. Systems should refer to the Guidance
Manual for Compliance with the Filtration and Disinfection Requirements for Public Works Systems
Using Surface Water Sources (AWWA, 1991) to ensure disinfectant schemes meet regulatory log
inactivation requirements. Utilities should also refer to EPA's Disinfection Profiling and
Benchmarking Guidance Manual (currently in production) to ensure compliance with the new
regulatory requirements of the IESWTR.
This chapter presents a brief discussion of the background and regulatory context of alternative
disinfectants, including an overview of the disinfection profiling and benchmarking approach to
evaluate disinfection efficiency. In addition, a decision-making framework is provided that utilities
can employ to assess the applicability of various disinfectants and disinfection strategies for
individual systems. Chapter 2 presents an overview of disinfection, including the use of chlorine,
with the next six chapters of this manual devoted to each of the following alternative disinfectants
and oxidants:
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• Chapter 3 - Ozone (63);
• Chapter 4 - Chlorine dioxide (C1O2);
• Chapter 5 - Potassium permanganate (KMnO4);
• Chapter 6 - Chloramine (NH2C1);
• Chapter 7 - Ozone/hydrogen peroxide combinations (Os/HbC^); and
• Chapter 8 - Ultraviolet radiation (UV).
For each disinfectant, this guidance manual describes the chemistry specific to the disinfection or
oxidation process, generation, primary uses and points of application, disinfection byproduct (DBF)
formation, pathogen inactivation and disinfection efficacy, the status of analytical methods for
residual monitoring, and operational considerations. Chapter 9 provides similar information
regarding the use of combined disinfectants. A summary of existing disinfectant usage in the United
States is provided in Appendix A. Cost estimates for the use of alternative disinfectants are provided
in Appendix B.
1.2 Background
The most important use of disinfectants in water treatment is to limit waterborne disease and
inactivate pathogenic organisms in water supplies. The first use of chlorine as a continuous process
in water treatment was in a small town in Belgium in the early 1900s (White, 1992). Since
introduction of filtration and disinfection at water treatment plants in the United States, waterborne
diseases such as typhoid and cholera have been virtually eliminated. For example, in Niagara Falls,
NY between 1911 and 1915, the number of typhoid cases dropped from 185 deaths per 100,000
population to nearly zero following introduction of filtration and chlorination (White, 1986).
In 1974, researchers in the Netherlands and the United States demonstrated that trihalomethanes
(THMs) are formed as a result of drinking water chlorination (Rook, 1974; Bellar et al., 1974).
THMs are formed when chlorine or bromide reacts with organic compounds in the water. EPA
subsequently conducted surveys confirming widespread occurrence of THMs in chlorinated water
supplies in the United States (Symons et al., 1975; USEPA, 1978). THMs and other DBPs have been
shown to be carcinogenic, mutagenic, etc. These health risks may be small, but with the large
population exposed, need to be taken seriously.
As a result of DBP concerns from chlorine, EPA, as well as the water treatment industry, placed
more emphasis on the use of disinfectants other than chlorine. Some of these alternative
disinfectants, however, have also been found to produce DBPs as a result of either reactions between
disinfectants and compounds in the water or as a natural decay product of the disinfectant itself
(McGuire et al., 1990; Legube et al; 1989). These DBPs include:
• Halogenated organics, such as THMs, haloacetic acids, haloketones, and others, that are
produced primarily as a result of chlorination.
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• Organic oxidation byproducts such as aldehydes, ketones, assimilable organic carbon (AOC), and
biodegradable organic carbon (BDOC), that are associated primarily with strong oxidants such as
ozone, chlorine, and advanced oxidation; and
• Inorganics such as chlorate and chlorite, associated with chlorine dioxide, and bromate, that is
associated with ozone, and has also has been found when chlorine dioxide is exposed to sunlight.
As documented in this manual, the type and amount of DBFs produced during treatment depends
largely on disinfectant type, water quality, treatment sequences, contact time, and environmental
factors such as temperature and pH.
When considering the use of alternative disinfectants, systems should ensure that the inactivation of
pathogenic organisms is not compromised. Pathogens pose an immediate critical public health threat
due to the risk of an acute disease outbreak. Although most identified public health risks associated
with DBFs are chronic, long-term risks, many systems will be able to lower DBF levels without
compromising microbial protection.
1.3 Regulatory Context
Pursuant to requirements of the Safe Drinking Water Act (SDWA) Amendments of 1996, EPA is
developing interrelated regulations to control microbial pathogens and disinfectant residuals and
disinfection byproducts in drinking water. These rules are collectively known as the
microbial/disinfection byproducts (M-DBP) rules and are intended to address complex risk trade-offs
between the desire to inactivate pathogens found in water and the need to reduce chemical
compounds formed as byproducts during disinfection. The rules are being promulgated in two
phases.
As part of the first phase, the Stage 1 DBPR and the IESWTR were promulgated on December 16,
1998 (63 FR 69390 and 63 FR 69478, respectively). The Stage 1 DBPR applies to all community
, water systems (CWS) and non-transient, non-community water systems (NTNCWS) that treat their
water with a disinfectant for either primary or residual treatment. The IESWTR amends the Surface
Water Treatment Rule (SWTR) and includes new and more stringent requirements to control
waterborne pathogens including specifically the protozoan Cryptosporidium. The IESWTR applies
to all public water systems that use surface water, or ground water under the direct influence of
surface water as defined at 40 CFR, Part 141, Subpart H1, and that serve at least 10,000 people.
Three future rules, also included in phase one include the Long-Term 1 Enhanced Surface Water
Treatment Rule (LT1ESWTR), the Ground Water Rule (GWR), and the Filter Backwash Rule
(FBR). Each of these rules are planned for promulgation in November 2000. The LT1ESWTR will
address pathogen inactivation and removal requirements for Subpart H systems serving fewer than
10,000 people. The GWR will specify appropriate disinfection techniques, including the use of best
1 Subpart H systems are defined as public water systems supplied by a surface water source or by a ground water source under the
direct influence of surface water.
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1. INTRODUCTION
management practices (BMPs) and source control measures. The FBR will set a standard for filter
backwash recycling for all public water systems regardless of size.
The second phase, consisting of the Stage 2 DBPR and the Long-Term 2 ESWTR, will be
promulgated in the year 2002 and will revisit the regulations for the formation of DBFs for all
systems and the inactivation and removal of pathogens for surface water systems, respectively.
The projected dates for future M-DBP regulatory activities are summarized in Table 1-1.
Table 1-1. Key Dates for Regulatory Activities
Date Regulatory Action
November 2000 Promulgate Long-Term 1 Enhanced Surface Water Treatment Rule
November 2000 Promulgate Ground Water Rule
November 2000 Promulgate Filter Backwash Rule
May 2002 Promulgate Stage 2 Disinfectants and Disinfection Byproduct Rule
May 2002 Promulgate Long-Term 2 Enhanced Surface Water Treatment Rule
Concurrent with the M-DBP rules, in May 1996, EPA promulgated the Information Collection Rule
(ICR) to obtain data on source water quality, byproduct formation, and drinking water treatment
plant design and operations. The ICR applies to Subpart H systems serving more than 100,000
people and ground water systems serving more than 50,000 people. EPA intended to use data from
the ICR to address completely the complex trade-offs between chronic DBP health risks and acute
pathogenic health risks, but delays in promulgation of the ICR eliminated this potential data source
for use in the IESWTR. Until the ICR data are analyzed in detail, EPA cannot fully address the issue
of DBP and pathogenic risk trade-offs.
National Primary and Secondary Drinking Water Regulations, published in 40 CFR Parts 141 and
143, respectively, limit the amount of specific contaminants and residual disinfectants and classes of
these compounds that are delivered to users of public water systems. These limits are expressed as
follows:
• Maximum Contaminant Level Goals (MCLGs). MCLGs are non-enforceable health goals for
public water systems. MCLGs are set at levels that, in the EPA Administrator's judgment, allow
no known or anticipated adverse effect on the health of persons to occur and that allow an
adequate margin of safety.
• Maximum Residual Disinfectant Level Goals (MRDLGs). As with MCLGs, EPA has
established MRDLGs for disinfectants at levels at which no known or anticipated adverse effects
on the health of persons occur and that allow an adequate margin of safety. MRDLGs are non-
enforceable health goals based only on health effects and exposure information and do not reflect
the benefit of the addition of the chemicals for control of waterborne microbial contaminants.
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• Maximum Contaminant Levels (MCLs). MCLs are enforceable standards set as close to the
MCLGs as technically and economically feasible.
• Maximum Residual Disinfectant Levels (MRDLs). MRDLs are similar to MCLs. MRDLs
are enforceable standards, analogous to MCLs, that recognize the benefits of adding a
disinfectant to water on a continuous basis and of addressing emergency situations such as
distribution system pipe ruptures. As with MCLs, EPA has set the MRDLs as close to the
MRDLGs as feasible.
In November 1979, the EPA set an interim MCL for Total THMs (TTHMs) of 0.10 mg/L as an
annual average for systems serving at least 10,000 people. This standard was based on the need to
reduce THM levels due to suspected carcinogenicity. Since then, EPA has developed and
promulgated standards for numerous contaminants. As of the December 16, 1998 DBPR
promulgation, MCLGs, MCLs, MRDLGs, and MRDLs are as presented in Tables 1-2 through 1-4.
As included in these tables, the December 16, 1998 Stage 1 DBPR:
• Lowered the existing MCL for TTHMs from 0.10 mg/L to 0.080 mg/L;
• Extended the MCL for TTHMs to all size systems;
• Requires enhanced coagulation or enhanced precipitative softening for certain systems;
• Established MRDLs and MRDLGs for chlorine, chloramine, and chlorine dioxide;
• Established MCLs for haloacetic acid (five) (HAAS), bromate, and chlorite, and
• Established MCLGs for eight disinfection byproducts.
Further, the SWTR of 1989 requires 3.0-log inactivation for Giardia cysts and 4.0-log inactivation
for viruses in surface water supplies. To meet these goals, the SWTR established treatment
requirements for filtration and disinfection. As shown in.Table 1-5, these goals can be met using
various treatment schemes that include filtration and disinfection. The JESWTR requires a 2.0 log
removal of Cryptosporidium for Subpart H systems serving at least 10,000 people.
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Table 1-2. Primary Drinking Water Regulations Related to Microbiological Contaminants
Compound
Gtardla lamblia
Legionella
Heterotrophic
Plate Count
Total Coliform
Turbidity
Viruses
MCLG
(mg/L)
Zero
Zero
N/A
Zero
N/A
Zero
MCL
(mg/L)
TT1
TT
TT
< 5.0%2
TT
TT
Potential Health Effects
Gastroenteric disease
Legionnaire's disease
Indicates water quality,
effectiveness of treatment
Indicates potential presence of
gastroenteric pathogens
Indicates water treatment failure
and pathogens in drinking water
Gastroenteric disease
Sources of Drinking Water
Contamination
Human and animal fecal waste
Common bacteria in natural
waters; can proliferate in water
heating systems
Human and animal fecal waste
Particles from storm runoff,
discharges into source water, and
erosion
Human and animal fecal waste
Source: AWWA Internet, 1997.
1TT = Treatment technique requirement in lieu of MCL as established in 40 CFR §141.70.
zNo more than 5.0 percent positive if >40 samples/month. No more than 1 positive if <40 samples/month [40 CFR §141.63(a) ].
Table 1-3. Primary Drinking Water Regulations Related to Disinfection Byproducts
Compound
Bromate
Bromodichloromethane
Brocnoform
Chlorite
Chloroform
Dibramochloromethane
Dlchloroacetic Acid
Haloacetic Acids1 (HAAS)
Trichloroacette Acid
Total Trihalomethanes2
(TTHMs)
MCLG
(mg/L)
Zero3
Zero3
Zero3
0.83
Zero3
0.063
Zero3
N/A
0.33
N/A
MCL
(mg/L)
0.010"
see TTHMs
see TTHMs
1.04
see TTHMs
see TTHMs
See HAAS
0.0604
See HAAS
0.084
Potential Health Effects
Cancer
Cancer, liver, kidney, and
reproductive effects
Cancer, nervous system,
liver and kidney effects
Hemolytic anemia
Cancer, liver, kidney,
reproductive effects
Nervous system, liver,
kidney, reproductive effects
Cancer and other effects
Cancer and other effects
Possibly cancer and
reproductive effects
Cancer and other effects
Sources of Drinking Water
Contamination
Ozonation byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water ozonation,
chloramination, and chlorination
byproduct
Chlorine dioxide disinfection
byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water chlorination and
chloramination byproduct
Drinking water chlorination and
chloramination byproduct
Source: 63 FR 69390 (12/16/98)
1 HAAS is the sum of the concentrations of mono-, di-, and trichloroacetic acids arid mono-and dibromoacetic acids.
2Total Trihalomethanes are the sum of the concentrations of bromodichlorbmethane, dibromochloromethane, bromoform, and chloroform.
3 Finalized on December 16,1998 (63 FR 69390) as established in 40 CFR §141.53.
* Finalized on December 16,1998 (63 FR 69390) as established in 40 CFR §141.64
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1. INTRODUCTION
Table 1-4. Primary Drinking Water Regulations Related to Residual Disinfectants
Disinfectant
Chlorine1
Chloramine2
Chlorine Dioxide
MRDLG3
(mg/L)
4 (as CI2)
4 (as CI2)
0.8(asCIO2) "
MRDL4
(mg/L)
4.0 (as CI2)
4.0 (as CI2)
0.8 (as CIO2)
Measured as free chlorine
2 Measured as total chlorine
3
Finalized on December 16,1998 (63 FR 69390) as established in 40 CFR §141.54.
Finalized on December 16,1998 (63 FR 69390) as established in 40 CFR §141.65
Table 1-5. Log Removal/Inactivation through Filtration and
Disinfection Required Under the SWTR
Process ._
Total log removalfmactivation Required
Conventional sedimentation/filtration credit
Disinfection inactivation required
Direct filtration credit
Disinfection inactivation required
Slow sand filtration credit
Disinfection inactivation required
Diatomaceous earth credit
Disinfection inactivation required
No Filtration
Disinfection inactivation required
Giardia cysts
3.0
2.5
' 0.5
2.0
1.0
2.0
1.0
2.0
1.0
0.0
3.0
Virus
4.0
2.0
2.0
1.0.
3.0
2.0
2.0
1.0
3.0
0.0
4.0
Source: AWWA, 1991.
Note: Some instances may require higher than 3 and 4 log removal. Also, some states may reduce removal filtration process.
1.3.1 Disinfection Profiling and Benchmarking
The IESWTR establishes disinfection benchmarking as a procedure requiring certain PWSs to
evaluate the impact on microbial risk of proposed changes in disinfection practice. It is designed to
facilitate utilities and States working together to assure that pathogen control is maintained while the
provisions of the Stage 1 DBPR are implemented. This procedure involves a PWS charting daily
levels of pathogen inactivation for a period of at least one year to create a profile of inactivation
performance. The PWS then uses this profile to determine a baseline or benchmark of inactivation
against which proposed changes in disinfection practices can be measured.
Systems are required to prepare a disinfection profile if either TTHM or HAA5 levels are at least
0.064 or 0.048 mg/L, respectively, as an annual average. These levels, equal to 80 percent of the
MCLs established for these compounds by the Stage 1 DBPR, are intended to include most systems
that will modify their disinfection practices to comply with the Stage 1 DBPR. To determine
applicability, systems that collected TTHM and HAAS data under the ICR must use the results of the
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last 12 months of ICR monitoring unless the State determines there is a more representative data set.
Non-ICR systems may use existing TTHM and HAAS data, if approved by the State, or must
conduct TTHM and HAAS monitoring for four quarters. This monitoring must be completed no later
than 15 months after promulgation of the IESWTR (i.e., by March, 2000). Alternatively, systems
can elect to forgo this monitoring if they construct a disinfection profile.
A disinfection profile consists of a compilation of daily Giardia lamblia log inactivations (plus virus
inactivations for systems using either chloramines or ozone for primary disinfection) computed over
a period of at least one year. It is based on daily measurements of disinfectant residual
concentration^), contact time(s), temperature, and pH. The profile may be developed using up to 3
years of existing (i.e. grandfathered) data, if the State finds the data acceptable. Systems having less
than 3 years of acceptable grandfathered data are required to conduct one year of monitoring to
create the profile. This monitoring must be complete within 27 months of IESWTR promulgation
(i.e., by March, 2001). The disinfection benchmark is equal to the lowest monthly average
inactivation level in the disinfection profile (or average of low months for multi-year profiles).
Any system required to develop a disinfection profile under the IESWTR that decides to make a
significant change to its disinfection practice must consult with the State prior to making the change.
Significant changes in disinfection practice are defined as: 1) moving the point of disinfection, not
including routine seasonal changes, 2) changing the type of disinfectant, 3) changing the disinfection
process, and 4) other modifications designated as significant by the State. As part of the consultation
process, the system must submit to the State the following information: a description of the proposed
change; the disinfection profile for Giardia lamblia (and, if necessary, viruses) and benchmark; and
an analysis of how the proposed change will affect the current levels of disinfection. In addition, the
State is required to review the disinfection profile a part of its periodic sanitary survey.
For more information on disinfection profiling and benchmarking, refer to EPA's Disinfection
Profiling and Benchmarking Guidance Manual (expected to be available in 1999).
1.4 Use of Disinfectants as Chemical Oxidants
Most disinfectants are strong oxidants and/or generate oxidants as byproducts (such as hydroxyl free
radicals) that react with organic and inorganic compounds in water. While the primary focus of this
manual is disinfection, many of the disinfectants described in this manual are also used for other
purposes in drinking water treatment, such as taste and odor control, improved flocculation, and
nuisance control. Because DBFs are produced irrespective of the intended purpose of the oxidant, it
is important to also address uses of disinfectants as oxidants in water treatment. These additional
uses are described in more detail in Chapter 2.
1.5 How Chlorine is Addressed in this Guidance Manual
This guidance manual does not provide as broad a discussion of chlorine and chlorination practices
as it does for the capabilities and uses of alternative disinfectants. There are two reasons EPA has
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1. INTRODUCTION
taken this approach: 1) the goal of this manual is to provide technical and engineering information to
utilities and the states on alternative disinfectants and oxidants, about which there is less
comprehensive information than for chlorine; and 2) a great majority of utilities already use chlorine,
in a wide variety of applications, for which there exists a wealth of literature on chlorine's uses and
performance capabilities. Summarizing this large body of knowledge in this guidance is neither
practical nor necessary.
This manual is not an EPA endorsement of alternative disinfectants nor is it a recommendation for
utilities to switch from chlorine to an alternative disinfectant. Rather, this manual provides technical
and engineering information to assist local professionals in making treatment decisions. EPA
believes that utility and state program personnel are best able to select disinfectants and design a
disinfection scheme, based upon site-specific conditions, that meets operational and regulatory
constraints. EPA does not require the use of chlorine or any other specific disinfectant for site-
specific uses. Again, local professionals are best suited to select disinfectants to address the unique
water treatment challenges posed by their source water and plant infrastructure.
EPA recognizes that, at the present time, chlorination is an important and central component of most
water treatment regimes in this country. As such, a summary of the uses and capabilities of chlorine
is provided in Section 2.7 of this manual. Section 2.7 also contains an extensive reference list for
additional information on chlorination.
1.6 A Summary of Alternative Disinfectant Properties
Subsequent chapters in this manual discuss several disinfectant alternatives available to a water
supplier. Table 1-6 summarizes the key technical and regulatory considerations associated with the
use of the various disinfectants for selecting the most appropriate disinfectant. The table provides
some broad guidelines to provide a framework for decision making. The ratings in Table 1-6 are
based on a typical disinfectant application. Thus, even though chlorine is considered to be prone to
THM formation, the table does not address the degree or amount of THMs produced. Similarly,
more than 2-log inactivation can be achieved for some disinfectants, but the high dose required may
not make it a reasonable application, and in that case, would be identified in the table as not able to
achieve 2-log inactivation.
The following key issues are addressed in Table 1-6:
• THMs, oxidized organics, and halogenated organics are produced. Halogenated organics are
formed when chlorine or ozone (in the presence bromide ion) is used, while oxidized organics
occur in the greatest concentration when a strong oxidant is used. The production of DBFs
depends on the amounts and types of precursors in the water.
• Inorganic byproducts are produced. Inorganic byproducts include chlorate ion, chlorite ion,
and bromate ion associated with chlorine dioxide and ozone.
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• MRDLs are required for some disinfectants. Note that this requirement must be balanced with
the requirement to maintain a residual in the distribution system. For most disinfectants, such as
chlorine, the MRDL is relatively high and will generally not create a problem.
• Lime softening impacts are noted. The high pH treatment during lime softening has an impact
on chlorine, chloramines, and UV.
• Turbidity impacts UV disinfection and ozonation. Ozone may interfere with coagulation and
settling. As such, it is recommended that ozonation be placed after settling but before filters to
minimize turbidity impacts.
• Inactivation requirements are divided into those achieving more or less than 2-log
inactivation. This differentiation is to identify the feasibility of achieving high inactivation at
modest doses. For example, while chlorine can achieve 3-log Giardio. cyst inactivation, the CT
requirement for 3-log inactivation of 100 to more than 300 mg-min/L will require high chlorine
doses and/or long contact times. However, 4-log virus inactivation is achievable with a CT of 15
to 60 mg-min/L for most temperatures.
Recently there has been some reports of 2-log and higher Cryptosporidium oocyst inactivation
with UV, using a system that concentrates the oocysts on a filter and then allows extended
exposure to UV to doses as high as 8,000 mW-s/cm2. This type application is considered
experimental at this time.
* Applicability as a secondary disinfectant represents the ability of the disinfectant to
maintain a residual in the distribution system. Only chlorine, chlorine dioxide, and
monochloramine provide residual disinfection in the distribution system. Chlorine dioxide is
limited to systems with smaller distribution systems because the total chlorine dioxide dose that
can be applied is limited by the production of chlorate ion and chlorite ion.
• Operator skill provides general guidance to the amount of operator attention and
maintenance required. All of the disinfectants can be placed on automatic control to limit the
amount of operator attention. The operational attention is rated based on the complexity of the
disinfectant application. Therefore, permanganate, which is a simple chemical feed system with
few mechanical elements, is rated 1 (low attention), while peroxone which include both ozone
systems and hydrogen peroxide feed systems, is rated 5 (high attention).
• All chemical disinfectants are judged to be applicable to small and large utilities. Modular
units of the technologies can cover a large range of flows. Ozone and chlorine dioxide generators
are available with small and large capacities. Chlorine and chemical feed systems have been
used successfully in all applications. Most UV water treatment facilities are less than 200 gpm in
capacity.
Key elements in the decision making process relate to the water source (i.e., ground or surface water)
and existing treatment configuration (including filtration) because these factors have a significant
impact on the degree of log removal required during disinfection. Water quality has a large impact on
the potential for DBF BOM formation.
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Table 1-6. Summary of Disinfectant Properties
(Based on Typical Disinfectant Application)
Condition
Produce THM with TOC
Produce oxidized organics
Produce halogenated organics
Produce inorganic byproducts ,
Produce BOM
MRDL applies
Lime softening impacts
Turbidity impacts
Meet Giardia - <2.0 log
Meet Giardia • >2.0 log
Meet Crypto • <2.0 log
Meet Crypto - >2.0 log
Meet Virus - <2.0 log
Meet Virus - >2.0 log
Secondary disinfectant
Operator skill (1=low; 5=high)
Applicable to large utilities
Applicable to small utilities
Chlorine
y
s
y
n
s
y
y
n
y
n
n
n
y
y
.y
1
y
y
«
C
o
(S
s
y
s
s
y
n
n
. s
y
y ,
y
y
y
y
n
1 5
y.
y
Chlorine Dioxide
n
s
n
y
s
y
n
n
y
y
y
n
y
y
s
5
y
y
Permanganate
n
s
n
n
n
n
n
n
n
n
n
n
n
n
n
1
y
y
Chloramine
y
n
y
n
n
y
y
n
n
n
n
n
n
n
y
2
y
y
Ozone/Peroxide
s
y
s
s
y
n
n
s
n
n
n
n
n
n
n
5
y
y
Ultraviolet
n
s
n
n
n
n
y
y
n
n
n
n
y
y
n
3
n
y
y = yes, n = no, s = sometimes
The following sections describe each of these phases.
1.7 Selecting a Disinfection Strategy
This section presents general guidance that can be used to assess the applicability of various
disinfectants or combination of disinfectants to select an appropriate disinfection strategy. Because
the selection of an appropriate strategy depends on site-specific conditions unique to each water
supply system, final selection of a strategy should be made with appropriate technical guidance (e.g.,
engineering study/evaluation or bench or pilot scale testing of alternatives). Selecting the most
appropriate disinfectant strategy for water treatment requires a balance among three key driving
forces:
• Providing water free of pathogens. Since the SWTR, the regulatory focus for pathogen
removal focused on coliform bacteria, heterotrophic plate counts, Giardia cysts, Legionella, and
viruses. Recently, the focus has been expanded to include Cryptosporidium oocyst removal and
inactivation, especially due to its resistance to chlorine.
April 1999
1-11
EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
1. INTRODUCTION
• Avoiding the production of disinfection byproducts (DBFs). Trihalomethanes (THMs), other
halogenated organics, ozone DBFs, oxidation byproducts, and some disinfectant residuals present
a health risk and must be limited in drinking waters. DBF precursor removal through process
optimization or enhanced coagulation is the first step in DBF control.
• Requiring residual disinfectant to maintain the bacteriological quality in the water as it is
distributed to customers to control regrowth. The potential for DBF formation increases with
extended contact between DBF precursors and residual disinfectants.
When changing disinfectants or oxidants water providers should consult with their primacy agencies.
The impact of the change should consider the impact on disinfection credit using disinfection
profiling and benchmarking techniques as summarized in Section 1.3.
1.7.1 Disinfection Strategy Evaluation
The selection of a disinfection strategy, as presented below, is affected by the following:
• Effectiveness of the current disinfection system;
• Need to change disinfectants;
• Selection of an alternative disinfectant; and
• Primary and secondary disinfection requirements.
As used in this guidance, primary and secondary disinfection are defined as follows:
Primary disinfection: The first (i.e., primary) disinfectant used in a treatment system, with the
primary objective of the disinfectant being to achieve the necessary CT (i.e., microbial
inactivation).
Secondary disinfection: The second disinfectant used in a treatment system, with the primary
objective of the disinfectant being to maintain the disinfection residual through the distribution
system.
For discussion, the approach to select a disinfection strategy is divided into three phases:
• Evaluate the current primary disinfection practice
• Select a primary disinfectant
• Select a secondary disinfectant.
1.7.1.1 Evaluate the Current Primary Disinfection Practice
Figure 1-1 presents the decision making process used to determine whether the present primary
disinfectant can meet disinfection and byproduct requirements. The key decision points in Figure 1-1
include:
• Meet current microbial inactivation limits. Microbial limits are defined by the primary
drinking water standards shown in Table 1-2. The regulated indicators of pathogens include
EPA Outdance Manual 1^12 April 1999
Alternative Disinfectants and Oxidants
-------
/. INTRODUCTION
Giardia lamblia, Legionella, HPC, total coliform turbidity, and viruses. The disinfectant must be
capable of meeting the inactivation requirements for disinfection. If not, the plant must
determine if the current disinfectant can meet the microbial inactivation requirements solely
through process modifications. A process modification may be to move the application point,
increase dose, increase contact time, or adjust pH. If not, a new disinfectant may be needed. In
some instances, a PWS may opt to improve its current microbial inactivation even though the
PWS is meeting current microbial limitations. In these instances, an evaluation is necessary to
ensure that any process modifications to improve inactivation will still provide compliance with
the microbial limits.
• Meet current DBF limits. A second set of limits imposed on disinfectant usage are DBF
requirements. The DBF limits are established in the Stage 1 DBF rule (See Table 1-3 and Table
1-5). To meet these limits on a consistent basis under normal varying water quality conditions,
80 percent of the MCL serves as an action level that requires a change in treatment practice.
Similar to microbial inactivation, some PWSs may desire to improve current DBF protection,
thus requiring evaluation of these modifications. By optimizing existing treatment processes, the
production of DBFs can be reduced. Optimization may include pretreatment optimization (i.e.,
coagulation, filtration, etc.) or process modifications such as moving the point of disinfection.
Enhanced coagulation is required by the Stage 1 DBF rule. Where process modifications are
contemplated, the PWS must ensure that both microbial inactivation and DBF formation comply
with applicable regulatory requirements. If optimized treatment cannot meet DBF and microbial
requirements, a new disinfectant may be needed.
1.7.1.2 Select a Primary Disinfectant
If it is determined that a new disinfectant is required or desired for better public health protection, the
second phase in the decision process (Figure 1 -2) addresses the factors concerning selection of a
primary disinfectant. This decision requires knowledge of the following three key components:
• TOC concentration. A high TOC concentration indicates a high potential for DBF formation.
In these cases, the decision tree will favor those disinfectants that will not produce DBFs or will
produce the least amount of DBFs. Note that precursor removal and enhanced coagulation are
used to reduce TOC during treatment optimization as indicated in Figure 1-1. "High TOC"
quantifies the potential to produce DBFs and is defined as a condition meeting one of the
following criteria:
- TOC exceeds 2 mg/L;
. - TTHM exceeds MCL (0.08 mg/L under Stage 1 DBPR); or
- HAAS exceeds MCL (0.06 mg/L under Stage 1 DBPR).
April 1999 1-13 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
1. INTRODUCTION
START j
Desire to
Improve
Current
Microbial
Inactivation?
Meet
Current
Microbial
Inactivation
Limits?
Change
Disinfectant
(See Rgure 1-2)
Meet Microbial
Goals with Process
Modification' ?
Desire to
Improve
Current DBF
Protection?
Meet
Current DBP
Limits?
Meet1
DBP Goals
with Process
Modifications?
Consult with Primacy
Agency when Changing
Treatment Process
Process Modification includes treatment optimization, moving the point of
dtsinl*ctant application, increasing disinfectant dosage, increasing contact time,
adjusting the pH> or other changes to existing treatment operations.
Figure 1-1. Flow Diagram to Evaluate Current Disinfection Practices
EPA Guidance Manual
Alternative Disinfectants and Oxidants
1-14
April 1999
-------
1. INTRODUCTION
• Bromide ion concentration. The reactions of strong oxidants (ozone and peroxone) with
bromide ion to produce hypobromous acid and bromate ion, precludes their usage with waters
containing high concentrations of bromide ion. High bromide ion is defined as concentrations
exceeding 0.10 mg/L.
• Filtered versus non-filtered systems. The use of ozone or ozone/peroxide for unfiltered
systems without the benefit of biofiltration to reduce ozonation byproducts and BOM is
strongly discouraged.
1.7.1.3 Select a Secondary Disinfectant
The selection of a secondary disinfectant depends on the selected primary disinfectant. Figure 1-3
identifies three decision points for secondary disinfectants:
• Assimilable organic carbon (AOC) concentration. AOC is produced when a strong oxidant
(e.g., ozone) is used as primary disinfectant in the presence of high TOC water. High AOC is
defined as concentrations exceeding 0.10 mg/L after filtration. In these cases, additional
biological or GAC treatment should be considered to stabilize the finished water and prevent
regrowth in the distribution system.
• DBF formation potential (DBPFP). The DBPFP serves as an indication of the amount of
organic byproducts that could be expected to form in the distribution system if chlorine is used.
Because DBF formation continues in the distribution system, the DBF content at the plant
effluent should be limited. A high DBPFP is defined as a water meeting one of the following
criteria:
- TTHM seven-day formation exceeds the MCL (0.08 mg/L under Stage 1 DBPR); or
- HAAS seven-day formation exceeds the MCL (0.06 mg/L under the Stage 1 DBPR).
• Distribution system retention time. In a large distribution system, booster stations may be
required to maintain the disinfection residual. Since chlorine dioxide has an upper limit for
application, its usage may not be feasible if relatively high doses are required to maintain a
residual in the distribution system. A distribution system retention time is considered high if it
exceeds 48 hours.
April 1999 1-15 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
1. INTRODUCTION
A. Subpart H Systems
that Filter
Chlorine Dioxide
Chlorine
Interactive Disinfectants
Need Bench/Pilot Study
Chlorine Dioxide
Chlorine
Ozone
Interactive Disinfectants
Yes
B. Subpart H Systems
that do not Filter
High Inactivation Reqmts.
High Disinfection Doses
Chlorine Dioxide
Chlorine
Interactive Disinfectants
Bench/Pilot Studies
Ozone
Chlorine Dioxide
Interactive Disinfectants
Excessive DBP Formation likely.
High disinfectant doses
required. May need TOC and
pathogen removal facilities.
Bench and pilot studies needed.
Consult Primacy Agency When
Changing Treatment Process
Yes
Excessive DBP Formation likely.
High disinfectant doses
required. May need TOC and
pathogen removal facilities.
Bench and pilot studies needed.
Figure 1-2. Flow Diagram to Narrow Selection of a New Primary Disinfectant
EPA Guidance Manual
Alternative Disinfectants and Oxidants
1-16
April 1999
-------
1. INTRODUCTION
Consult Primacy Agency When
Changing Treatment Process
Chlorine
Chloramine
Chlorine
Chlorine Dioxide
Chloramine
Chlorine Dioxide
Chloramine
Chloramine
Figure 1-3. Flow Diagram to Narrow Selection of a New Secondary Disinfectant
1.7.2 Summary
The approach outlined above serves as a basis to evaluate the need for and to select the most
appropriate alternative disinfectants for drinking water systems. The approach is general enough to
cover the likely outcomes of the inactivation requirements and DBF formation. However, in some
cases, site-specific conditions may dictate a different approach. It is important to consult with the
primacy agency whenever a change in treatment is considered. Remember that the IESWTR requires
certain PWSs to use disinfection profiling and benchmarking procedures when proposing changes to
disinfection practices. In addition, some of the decisions may lead to a situation where additional
treatment will be required. For example, filtration may be needed to reduce the disinfectant dose and
April 1999
1-17
EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
7. INTRODUCTION
limit DBF formation in cases of high TOC and high bromide ion levels. In those instances, bench
scale or pilot studies may be required to select the most appropriate disinfectant.
1.8 References
I. AWWA (American Water Works Association). 1991. Guidance Manual for Compliance with
the Filtration and Disinfection Requirements for Public Works Systems Using Surface Water
Sources.
2. AWWA Safe Drinking Water Advisor - Library on Internet (1997).
3. Bellar, T.A, JJ. Lichtenberg, and R.C. Kroner. 1974. "The Occurrence of Organohalides in
Chlorinated Drinking Water." /. AWWA. 66(12): 703-706.
4. Legube, B., J.P. Croue', J. De Latt, and M. Dore'. 1989. "Ozonation of an Extracted Aquatic
Fulvic Acid: Theoretical and Practical Aspects." Ozone Sci. Eng. 11(1): 69-91.
5. McGuire, M.J., D.W. Ferguson, and J.T. Gramith. 1990. Overview of Ozone Technology for
Organics Control and Disinfection. Conference proceedings, AWWA Seminar on Practical
Experiences with Ozone for Organics Control and Disinfection, Cincinnati, OH.
6. Rook, JJ. 1974. "Formation of Haloforms during Chlorination of Natural Water." Water
Treatment and Examination. 23(2): 234-243.
7. Sawyer, C.N., P.L. McCarty, L. Parkin, and G.F. Parkin. 1994. Chemistry for Environmental
Engineering, fourth edition. McGraw Hill, Inc., New York, NY.
8. Symons, J.M., T.A. Bellar, J.K. Carswell, J. DeMarco, K.L. Kropp, G.G. Robeck, D.R. Seeger,
C.J. Slocum, B.L. Smith and A.A. Stevens. 1975. "National Organics Reconnaissance Survey
for Halogenated Organics." J. AWWA. 67(11): 634-647.
9. USEPA (U.S. Environmental Protection Agency). 1997. Community Water System Survey -
Volume II. Detailed Survey Result Tables and Methodology Report. EPA 815-R-97-001b.
10. USEPA. 1978. National Organics Monitoring Survey (NOMS). Technical Support Division,
U.S. Environmental Protection Agency, Office of Drinking Water. Cincinnati, OH.
11. White, G.C. 1992. Handbook of Chlorination and Alternative Disinfectants. Vol. 3. Van
Nostrand Reinhold Co. New York, NY.
12. White, G.C. 1986. Handbook of Chlorination. Van Nostrand Reinhold Company, New York,
NY.
EPA Guidance Manual 1-18 April 1999
Alternative Disinfectants and Oxidants
-------
2. DISINFECTANT USE IN WATER
TREATMENT
To comply with the SDWA regulations, the majority of PWSs use some form of water treatment.
The 1995 Community Water Systems Survey (USEPA, 1997a) reports that in the United States, 99
percent of surface water systems provide some treatment to their water, with 99 percent of these
treatment systems using disinfection/oxidation as part of the treatment process. Although 45 percent
of ground water systems provide no treatment, 92 percent of those ground water plants that do
provide some form of treatment include disinfection/oxidation as part of the treatment process. The
most commonly used disinfectants/oxidants (in no particular order) are chlorine, chlorine dioxide,
chloramines, ozone, and potassium permanganate.
Disinfectants are also used to achieve other specific objectives in drinking water treatment. These
other objectives include nuisance control (e.g., for zebra mussels and Asiatic clams), oxidation of
specific compounds (i.e., taste and odor causing compounds, iron, and manganese), and use as a
coagulant and filtration aid.
The purpose of this chapter is to:
• Provide a brief overview of the need for disinfection in water treatment.
• Provide basic information that is common to all disinfectants;
• Discuss other uses for disinfectant chemicals (i.e., as oxidants);
• Describe trends in DBP formation and the health effects of DBFs found in water treatment;
• Discuss microorganisms of concern in water systems, their associated health impact, and the
inactivation mechanisms and efficiencies of various disinfectants; and
• Summarize current disinfection practices in the United States, including the use of chlorine as a
disinfectant and an oxidant.
2.1 Need for Disinfection in Water Treatment
Although the epidemiological relation between water and disease had been suggested as early as the
1850s, it was not until the establishment of the germ theory of disease by Pasteur in the mid-1880s
that water as a carrier of disease-producing organisms was understood. In the 1880s, while London
experienced the "Broad Street Well" cholera epidemic, Dr. John Snow conducted his now famous
epidemiological study. Dr. Snow concluded that the well had become contaminated by a visitor,
with the disease, who had arrived in the vicinity. Cholera was one of the first diseases to be
Anrii HQQQ; 2^\ EPA Guidance Manual
Mp Alternative Disinfectants and Oxidants
-------
2. DISINFECTANT USE IN WATER TREATMENT
recognized as capable of being waterborne. Also, this incident was probably the first reported
disease epidemic attributed to direct recycling of non-disinfected water. Now, over 100 years later,
the list of potential waterborne diseases due to pathogens is considerably larger, and includes
bacterial, viral, and parasitic microorganisms, as shown in Table 2-1, Table 2-2 and Table 2-3,
respectively.
A major cause for the number of disease outbreaks in potable water is contamination of the
distribution system from cross connections and back siphonage with non-potable water. However,
outbreaks resulting from distribution system contamination are usually quickly contained and result
in relatively few illnesses compared to contamination of the source water or a breakdown in the
treatment system, which typically produce many cases of illnesses per incident. When considering
the number of cases, the major causes of disease outbreaks are source water contamination and
treatment deficiencies (White, 1992). For example, in 1993 a Cryptosporidiosis outbreak affected
over 400,000 people in Milwaukee, Wisconsin. The outbreak was associated with deterioration in
the raw water quality and a simultaneous decrease in the effectiveness of the coagulation-filtration
process (Kramer etal., 1996; MacKenzie et al., 1994). Historically, about 46 percent of the
outbreaks in the public water systems are found to be related to deficiencies in source water and
treatment systems with 92 percent of the causes of illness due to these two particular problems.
All natural waters support biological communities. Because some microorganisms can be
responsible for public health problems, biological characteristics of the source water are one of the
most important parameters in water treatment. In addition to public health problems, microbiology
can also affect the physical and chemical water quality and treatment plant operation.
2.1.1 Pathogens of Primary Concern
Table 2-4 shows the attributes of three groups of pathogens of concern in water treatment, namely
bacteria, viruses, and protozoa.
2.1.1.1 Bacteria
Bacteria are single-celled organisms typically ranging in size from O.I to 10 jam. Shape,
components, size, and the manner in which they grow can characterize the physical structure of the
bacterial cell. Most bacteria can be grouped by shape into four general categories: spheroid, rod,
curved rod or spiral, and filamentous. Cocci, or spherical bacteria, are approximately 1 to 3 urn in
diameter. Bacilli (rod-shaped bacteria) are variable in size and range from 0.3 to 1.5 um in width (or
diameter) and from 1.0 to 10.0 jam in length. Vibrios, or curved rod-shaped bacteria, typically vary
in size from 0.6 to 1.0 urn in width (or diameter) and from 2 to 6 um in length. Spirilla (spiral
bacteria) can be found in lengths up to 50 um whereas filamentous bacteria can occur in length in
excess of 100 urn.
EPA Guidance Manual 2-2 April 1999
Alternative Disinfectants and Oxidants
-------
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2. DISINFECTANT USE IN WATER TREATMENT
Table 2-3. Waterborne Diseases from Parasites
Causative Agent
Disease
Symptoms
Ascario
lumbricoides
(round worm)
Ciyptosporidium
muris and parvum
Entamoeba
histolytlca
Glardia lamblla
Naegleria gruberi
Schistosoma
manson!
Taenia saginata
(beef tapeworm)
Ascariasis
Vomiting, live worms in feces
Cryptosporidiosis
Amebiasis
Acute diarrhea, abdominal pain, vomiting, and low-grade fever.
Can be life-threatening in immunodeficient patients
Diarrhea alternating with constipation, chronic dysentery with
mucus and blood
Giardiasis
Amoebic
meningoecephalitis
Schistosomiasis
Intermittent diarrhea
Death
Liver and bladder infection
Taeniasis
Abdominal pain, digestive disturbances, loss of weight
Source: Geldreich, 1972; Beneson, 1981.
2.1.1.2 Viruses
Viruses are microorganisms composed of the genetic material deoxyribonucleic acid (DNA) or
ribonucleic acid (RNA) and a protective protein coat (either single, double, or partially double
stranded). All viruses are obligate parasites, unable to carry out any form of metabolism and are
completely dependent upon host cells for replication. Viruses are typically 0.01 to 0.1 um in size and
are very species specific with respect to infection, typically attacking only one type of host.
Although the principal modes of transmission for the hepatitis B virus and poliovirus are through
food, personal contact, or exchange of body fluids, these viruses can be transmitted through potable
water. Some viruses, such as the retroviruses (including the HIV group), appear to be too fragile for
water transmission to be a significant danger to public health (Riggs, 1989).
2.1.1.3 Protozoa
Protozoa are single-cell eucaryotic microorganisms without cell walls that utilize bacteria and other
organisms for food. Most protozoa are free-living in nature and can be encountered in water;
however, several species are parasitic and live on or in host organisms. Host organisms can vary
from primitive organisms such as algae to highly complex organisms such as human beings. Several
species of protozoa known to utilize human beings as hosts are shown in Table 2-5.
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Table 2-4. Attributes of the Three Waterborhe Pathogens of Concern in Water
Treatment
Organism , IZ^ Mobility Point(s) of Origin
(Hm)
Bacteria 0.1-10 Motile, Humans and
Nonmotile animals, water, and
contaminated food
Viruses 0.01-0.1 Nonmotile Humans and
animals, polluted
water, and
contaminated food
Protozoa 1-20 Motile, Humans and
Nonmotile animals, sewage,
decaying vegetation,
and water
Resistance to
Disinfection
Type specific -
bacterial spores
typically have the
highest resistance
whereas vegetative
bacteria have the
lowest resistance
Generally more
resistant than
vegetative bacteria
More resistant than
viruses or vegetative
bacteria
Removal by
Sedimentation,
Coagulation, and
Filtration
Good, 2 to 3-log
removal
Poor, 1 to 3-log
removal
Good, 2 to 3-log
removal
Table 2-5. Human Parasitic Protozoans
Protozoan
Host(s)
Disease
Transmission
Occurrence
Acanathamoeba Fresh water,
castellannii sewage, humans,
and soil
Balantidium coll Pigs, humans
Cryptosporidium Animals, humans
parvum
Entamoeba Humans
histolytica
Giardia lamblia Animals, humans
Naegleria fowleri Soil, water,
humans and
decaying
vegetation
Amoebic
meningoencephalitis
Balantidiasis
(dysentery)
Cryptosporidiosis
Gains entry through
abrasions, ulcers, and as
secondary invader during
other infections
Contaminated water
Person-to-person or
animal-to-person contact,
ingestion of fecally
contaminated water or
food, or contact with
fecally contaminated
environmental surfaces.
North America
Micronesia has
been the only known
site of an outbreak
Canada, England,
and the United
States
Amoebic dysentery Contaminated water
Giardiasis
(gastroenteritis)
Primary amoebic
meningoencephalitis
Contaminated water
Nasal inhalation with
subsequent penetration
of nasopharynx;
exposure from swimming
in freshwater lakes
Last United States
outbreak, 1953
Mexico, United
States, USSR
North America
Source: Montgomery, 1985; AWWA, 1995.
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2.1.2 Recent Waterborne Outbreaks
Within the past 40 years, several pathogenic agents never before associated with documented
waterborne outbreaks have appeared in the United States. Enteropathogenic E. coli and Giardia
lamblia were first identified to be the etiological agent responsible for waterborne outbreaks in the
1960s. The first recorded Cryptosporidium infection in humans occurred in the mid-1970s. Also
during that time was the first recorded outbreak of pneumonia caused by Legionella pneumophila
(Centers for Disease Control, 1989; Witherell et al., 1988). Recently, there have been numerous
documented waterborne disease outbreaks that have been caused by E. coli, Giardia lamblia,
Cryptosporidium, and Legionella pneumophila.
2.1.2.1 E. coli
The first documented case of waterborne disease outbreaks in the United States associated with
enteropathogenic E. coli occurred in the 1960s. Various serotypes of E. coli have been implicated as
the etiological agent responsible for disease in newborn infants, usually the result of cross
contamination in nurseries. Now, there have been several well-documented outbreaks of E. coli
(serotypes 0111:B4 and 0124:B27) associated with adult waterborne disease (AWWA, 1990, and
Craun, 1981). In 1975, the etiologic agent of a large outbreak at Crater Lake National Park was E.
coli serotype 06:H16 (Craun, 1981).
2.1.2.2 Giardia lamblia
Similar to E. coli, Giardia lamblia was first identified in the 1960s to be associated with waterborne
outbreaks in the United States. Giardia lamblia is a flagellated protozoan that is responsible for
Giardiasis, a disease that can range from being mildly to extremely debilitating. Giardia is .currently
one of the most commonly identified pathogens responsible for waterborne disease outbreaks. The
life cycle of Giardia includes a cyst stage when the organism remains dormant and is extremely
resilient (i.e., the cyst can survive some extreme environmental conditions). Once ingested by a
warm-blooded animal, the life cycle of Giardia continues with excystation. The cysts are relatively
large (8-14 jam) and can be removed effectively by filtration using diatomaceous earth, granular
media, or membranes.
Giardiasis can be acquired by ingesting viable cysts from food or water or by direct contact with
fecal material. In addition to humans, wild and domestic animals have been implicated as hosts.
Between 1972 and 1981, 50 waterborne outbreaks of Giardiasis occurred with about 20,000 reported
cases (Craun and Jakubowski, 1986). Currently, no simple and reliable method exists to assay
Giardia cysts in water samples. Microscopic methods for detection and enumeration are tedious and
require examiner skill and patience. Giardia cysts are relatively resistant to chlorine, especially at
higher pH and low temperatures.
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2.1.2.3 Cryptosporidium
Cryptosporidium is a protozoan similar to Giardia. It forms resilient oocysts as part of its life cycle.
The oocysts are smaller than Giardia cysts, typically about 4-6 um in diameter. These oocysts can
survive under adverse conditions until ingested by a warm-blooded animal and then continue with
excy station.
Due to the increase in the number of outbreaks of Cryptosporidiosis, a tremendous amount of
research has focused on Cryptosporidium within the last 10 years. Medical interest has increased
because of its occurrence as a life-threatening infection to individuals with depressed immune
systems. As previously mentioned, in 1993, the largest documented waterborne disease outbreak in
United States history occurred in Milwaukee and was determined to be caused by Cryptosporidium.
An estimated 403,000 people became ill, 4,400 people were hospitalized, and 100 people died. The
outbreak was associated with a deterioration in raw water quality and a simultaneous decrease in
effectiveness of the coagulation-filtration process, which led to an increase in the turbidity of treated
water and inadequate removal of Cryptosporidium oocysts.
2.1.2.4 Legionella pneumophila
An outbreak of pneumonia occurred in 1976 at the annual convention of the Pennsylvania American
Legion. A total of 221 people were affected by the outbreak, and 35 of those afflicted died. The
cause of the pneumonia was not determined immediately despite an intense investigation by the
Centers for Disease Control. Six months after the incident, microbiologists were able to isolate a
bacterium from the autopsy lung tissue of one of the Legionnaires. The bacterium responsible to the
outbreak was found to be distinct from other known bacterium and was named Legionella
pneumophila (Witherell et al., 1988). Following the discovery of this organism, other Legionella-
like organisms were discovered. Altogether, 26 species of Legionella have been identified, and
seven are etiologic agents for Legionnaires' disease (AWWA, 1990).
Legionnaires' disease does not appear to be transferred person-to-person. Epidemiological studies
have shown that the disease enters the body through the respiratory system. Legionella can be
inhaled in water particles less than Sum in size from facilities such as cooling towers, hospital hot
water systems, and recreational whirlpools (Witherell et al., 1988).
2.1.3 Mechanism of Pathogen Inactivation
The three primary mechanisms of pathogen inactivation are to:
• Destroy or impair cellular structural organization by attacking major cell constituents, such as
destroying the cell wall or impairing the functions of semi-permeable membranes;
• Interfere with energy-yielding metabolism through enzyme substrates in combination with
prosthetic groups of enzymes, thus rendering the enzymes non-functional; and
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• Interfere with biosynthesis and growth by preventing synthesis of normal proteins, nucleic acids,
coenzymes, or the cell wall.
Depending on the disinfectant and microorganism type, combinations of these mechanisms can also
be responsible for pathogen inactivation. In water treatment, it is believed that the primary factors
controlling disinfection efficiency are: (1) the ability of the disinfectant to oxidize or rupture the cell
wall; and (2) the ability of the disinfectant to diffuse into the cell and interfere with cellular activity
(Montgomery, 1985).
2.2 Other Uses of Disinfectants in Water Treatment
Disinfectants are used for more than just disinfection in drinking water treatment. While inactivation
of pathogenic organisms is a primary function, disinfectants are also used oxidants in drinking water
treatment for several other functions:
• Minimization of DBF formation;
• Control of nuisance Asiatic clams and zebra mussels;
• Oxidation of iron and manganese;
• Prevention of regrowth in the distribution system and maintenance of biological stability;
• Removal of taste and odors through chemical oxidation;
• Improvement of coagulation and filtration efficiency;
• Prevention of algal growth in sedimentation basins and filters;
• Removal of color.
A brief discussion of these additional oxidant uses follows.
2.2.1 Minimization of DBF Formation
Strong oxidants may play a role in disinfection and DBP control strategies in water treatment.
Several strong oxidants, including potassium permanganate and ozone, may be used to control DBP
precursors.
Potassium permanganate can be used to oxidize organic precursors at the head of the treatment plant,
thus minimizing the formation of byproducts at the downstream disinfection stage of the plant. The
use of potassium permanganate as an oxidant and disinfectant is discussed in Chapter 5 of this
guidance manual.
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The use of ozone for oxidation of DBF precursors is currently being studied. Early work has
shown that the effects of ozonation, prior to chlorination, were highly site-specific and
unpredictable. The key variables that seem to determine the effect of ozone are dose, pH,
alkalinity, and the nature of the organic material. Ozone has been shown to be effective for DBF
precursor reduction at low pHs. However, at higher pHs (i.e., above 7.5), ozone may actually
increase the amount of chlorination byproduct precursors. The use of ozone as an oxidant and
disinfectant is addressed in detail in Chapter 3 of this document.
2.2.2 Control of Nuisance Asiatic Clams and Zebra Mussels
The Asiatic clam (Corbiculafluminea) was introduced to the United States from Southeast Asia in
1938 and now inhabits almost every major river system south of 40° latitude (Britton and Morton,
1982; Counts, 1986). Asiatic clams have been found in the Trinity River, TX; the Ohio River at
Evansville, IN; New River at Narrows and Glen Lyn, VA; and the Catawba River in Rock Hill, SC
(Belanger et al., 1991; Cameron et al., 1989a; Matisoff et al., 1996). This animal has invaded many
water utilities, clogging source water transmission systems and valves, screens, and meters;
damaging centrifugal pumps; and causing taste and odor problems (Sinclair, 1964; Evans et al., 1979;
Smith, 1979).
Cameron et al. (1989a) investigated the effectiveness of several oxidants to control the Asiatic clam
in both the juvenile and adult phases. As expected, the adult clam was found to be much more
resistant to oxidants than the juvenile form. In many cases, the traditional method of control, free
chlorination, cannot be used because of the formation of excessive amounts of THMs. As shown in
Table 2-6, Cameron et al. (1989a) compared the effectiveness of four oxidants for controlling the
juvenile Asiatic clam in terms of the LT50 (time required for 50 percent mortality).
Monochloramine was found to be the best for controlling the juvenile clams without forming THMs.
Further research showed that the effectiveness of monochloramine increased greatly as the
temperature increased (Cameron et al., 1989b). Note that the temperatures in this study reflect
conditions in the Lynchburg Reservoir, Houston, Texas. Clams can tolerate temperatures between 2
and 35°C (Cameron et al. 1989a).
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Table 2-6. The Effects of Various Oxidants on Mortality of the
Asiatic Clam (Corbicula fluminea)
Chemical R,esid"al
(mg/L)
Free chlorine
Potassium
Permanganate
Monochloramine
Chlorine dioxide
0.5
4.8
4.7
1.1
4.8
2.6
10.7
1.2
4.7
Temperature
fC) PH
23
21
16
17
17
28
17
24
22
8.0
7.9
7.8
7.6
7.6
7.9
7.9
6.9
6.6
Life Stage
A
A
J
J
J
J
J
J
J
LT50
(days)
8.7
5.9
4.8
7.9
8.6
0.6
0.5
0.7
0.6
AsAdult; J=Juvenile
Source: Cameron etal., 1989a.
In a similar study, Belanger et al. (1991) studied the biocidal potential of total residual chlorine,
monochloramine, monochloramine plus excess ammonia, bromine, and copper for controlling the
Asiatic clam. Belanger et al. (1991) showed that monochloramine with excess ammonia was the
most effective for controlling the clams at 30°C. Chlorination at 0.25 to 0.40 mg/L total residual
chlorine at 20 to 25°C controlled clams of all sizes (LT50 below 28 days) but had minimal effect at
12 to 15°C (as low as zero mortality). As in other studies, the toxicity of all the biocides was highly
dependent on temperature and clam size.
The zebra mussel (Dreissena polymorphd) is a recent addition to the fauna of the Great Lakes. It was
first found in Lake St. Glair in 1988, though it is believed that this native of the Black and Caspian
seas, was brought over from Europe in ballast water around 1985 (Herbert et al., 1989). The zebra
mussel population in the Great Lakes has expanded very rapidly, both in size and geographical
distribution (Roberts, 1990). Lang (1994) reported that zebra mussels have been found in the Ohio
River, Cumberland River, Arkansas River, Tennessee River, and the Mississippi River south to New
Orleans.
Klerks and Fraleigh (1991) evaluated the effectiveness of hypochlorite, permanganate, and hydrogen
peroxide with iron for their effectiveness controlling adult zebra mussels. Both continuous and
intermittent 28-day static renewal tests were conducted to determine the impact of intermittent
dosing. Intermittent treatment proved to be much less effective than continuous dosing.
The hydrogen peroxide-iron combination (1-5 mg/L with 25 percent iron) was less effective in
controlling the zebra mussel than either permanganate or hypochlorite. Permanganate (0.5-2.5 mg
KMnO4/L) was usually less effective than hypochlorite (0.5-10 mg C12/L).
Van Benschoten et al. (1995) developed a kinetic model to predict the rate of mortality of the zebra
mussel in response to chlorine. The model shows the relationship between chlorine residual and
temperature on the exposure time required to achieve 50 and 95 percent mortality. Data were
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collected for chlorine residuals between 0.5 and 3.0 mg C12/L and temperatures from 0.3 to 24°C.
The results show a strong dependence on temperature and required contact times ranging from two
days to more than'a month, depending on environmental factors and mortality required.
Brady et al. (1996) compared the efficiency of chlorine to control growth of zebra mussel and quagga
mussel (Dreissena bugensis). The quagga mussel is a newly identified mollusk within the Great
Lakes that is similar in appearance to the zebra mussel. Full-scale chlorination treatment found a
significantly higher mortality for the quagga mussel. The required contact time for 100 percent
mortality for quagga and zebra mussels was 23 days and 37 days, respectively, suggesting that
chlorination programs designed to control zebra mussels should also be effective for controlling
populations of quagga mussels.
Matisoff et al. (1996) evaluated chlorine dioxide (C1O2) to control adult zebra mussels using single,
intermittent, and continuous exposures. A single 30-minute exposure to 20 mg/L chlorine dioxide or
higher concentration induced at least 50 percent mortality, while sodium hypochlorite produced only
26 percent mortality, and permanganate and hydrogen peroxide were totally ineffective when dosed
at 30 mg/L for 30 minutes under the same conditions. These high dosages, even though only used
for a short period, may not allow application directly in water for certain applications due to
byproducts that remain in the water. Continuous exposure to chlorine dioxide for four days was
effective at concentrations above 0.5 mg/L (LC50 = 0.35 mg/L), and 100 percent mortality was
achieved at chlorine dioxide concentrations above 1 mg/L.
These experiences all show that the dose required to induce mortality to these nuisance organisms is
extremely high, both in terms of chemical dose and contact time. The potential impact on DBFs is
significant, especially when the water is high in organic content with a high propensity to form
THMs and other DBFs.
2.2.3 Oxidation of Iron and Manganese
Iron and manganese occur frequently in ground waters but are less problematic in surface waters.
Although not harmful to human health at the low concentrations typically found in water, these
compounds can cause staining and taste problems. These compounds are readily treated by oxidation
to produce a precipitant that is removed in subsequent sedimentation and filtration processes.
Almost all the common oxidants except chloramines will convert ferrous (2+) iron to the ferric (3+)
state and manganese (2+) to the (4+) state, which will precipitate as ferric hydroxide and manganese
dioxide, respectively (AWWA, 1990). The precise chemical composition of the precipitate will
depend on the nature of the water, temperature, and pH.
Table 2-7 shows that oxidant doses for iron and manganese control are relatively low. In addition,
the reactions are relatively rapid, on the order of seconds while DBF formation occurs over hours.
Therefore, with proper dosing, residual chlorine during iron and manganese oxidation is therefore
relatively low and short lived. These factors reduce the potential for DBF formation as a result of
oxidation for iron and manganese removal.
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Table 2-7. Oxidant Doses Required for Oxidation of Iron and Manganese
ll1 " " ' I
Oxidant
Chlorine, CI2
Chlorine dioxide, ClOa
Ozone, Oa
Oxygen, O2
Potassium permanganate, KMnO4
Iron (II)
(mg/mg Fe)
0.62
1.21
0.43
0.14
0.94
Manganese (II)
(mg/mg Mn)
0.77
2.45
0.88*
0.29
1.92
Source: Culp/Wesner/Culp, 1986; Langlais etal., 1991.
* Optimum pH for manganese oxidation using ozone is 8-8.5 Source: Reckhow et al., 1991.
2.2.4 Prevention of Regrowth in the Distribution System and
Maintenance of Biological Stability
Biodegradable organic compounds and ammonia in treated water can cause microbial growth in the
distribution system. "Biological stability" refers to a condition wherein the treated water quality does
not enhance biological growth in the distribution system. Biological stability can be accomplished in
several ways:
• Removing nutrients from the water prior to distribution;
• Maintaining a disinfectant residual in the treated water; and
• Combining nutrient removal and disinfectant residual maintenance.
To maintain biological stability in the distribution system, the Total Coliform Rule (TCR) requires
that treated water have a residual disinfectant of 0.2 mg/L when entering the distribution system. A
measurable disinfectant residual must be maintained in the distribution system, or the utility must
show through monitoring that the heterotrophic plate count (HPC) remains less than 500/100 mL. A
system remains in compliance as long as 95 percent of samples meet these criteria. Chlorine,
monochloramine, and chlorine dioxide are typically used to maintain a disinfectant residual in the
distribution system. Filtration can also be used to enhance biological stability by reducing the
nutrients in the treated water.
The level of secondary disinfectant residual maintained is low, typically in the range of 0.1 -0.3 mg/L,
depending on the distribution system and water quality. However, because the contact times in the
system are quite long, it is possible to generate significant amounts of DBFs in the distribution
system, even at low disinfectant doses.
Distribution system problems associated with the use of combined chlorine residual (chloramines), or
no residual, have been documented in several instances. The use of combined chlorine is
characterized by an initial satisfactory phase in which chloramine residuals are easily maintained
throughout the system and bacterial counts are very low. However, problems may develop over a
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period of years including increased bacterial counts, reduced combined chlorine residual, increased
taste and odor complaints, and reduced transmission main carrying capacity. Conversion of the
system to free-chlorine residual produces an initial increase in consumer complaints of taste and
odors resulting from oxidation of accumulated organic material. Also, it is difficult to maintain a
free-chlorine concentration at the ends of the distribution system (AWWA, 1990).
2.2.5 Removal of Taste and Odors Through Chemical Oxidation
Tastes and odors in drinking water are caused by several sources, including microorganisms,
decaying vegetation, hydrogen sulfide, and specific compounds of municipal, industrial, or
agricultural origin. Disinfectants themselves can also create taste and odor problems. In addition to
a specific taste-and odor-causing compound, the sensory impact is often accentuated by a
combination of compounds. More recently, significant attention has been given to tastes and odors
from specific compounds such as geosmin, 2-methylisoborneol (MIB), and chlorinated inorganic and
organic compounds (AWWARF, 1987).
Oxidation is commonly used to remove taste and odor causing compounds. Because many of these
compounds are very resistant to oxidation, advanced oxidation processes (ozone/hydrogen peroxide,
ozone/UV, etc.) and ozone by itself are often used to address taste and odor problems. The
effectiveness of various chemicals to control taste and odors can be site-specific. Suffet et al. (1986)
found that ozone is generally the most effective oxidant for use in taste and odor treatment. They
found ozone doses of 2.5 to 2.7 mg/L and 10 minutes of contact time (residual 0.2 mg/L)
significantly reduce levels of taste and odors. Lalezary et al. (1986) used chlorine, chlorine dioxide,
ozone, and permanganate to treat earthy-musty smelling compounds. In that study, chlorine dioxide
was found most effective, although none of the oxidants were able to remove geosmin and MIB by
more than 40 to 60 percent. Potassium permanganate has been used in doses of 0.25 to 20 mg/L.
Studies at the Metropolitan Water District of Southern California demonstrated the effectiveness of
peroxone (ozone plus hydrogen peroxide) to remove geosmin and MIB in water treatment (Ferguson
et al., 1990; Ferguson et al., 1991; Huck et al., 1995).
Prior experiences with taste and odor treatment indicate that oxidant doses are dependent on the
source of the water and causative compounds. In general, small doses can be effective for many taste
and odor compounds, but some of the difficult-to-treat compounds require strong oxidants such as
ozone and/or advanced oxidation processes or alternative technologies such as granular activated
carbon (GAC) adsorption.
2.2.6 Improvement of Coagulation and Filtration Efficiency
Oxidants, specifically ozone, have been reported to improve coagulation and filtration efficiency (Gurol
andPidatella, 1983; Farvardin and Collins, 1990; Reckhow et al., 1993; Masschelein, 1992). Others,
however, have found no improvement in effluent turbidity from oxidation (Tobiason et al., 1992;
Hiltebrand et al., 1986). Prendiville (1986) collected data from a large treatment plant showing that
preozonation was more effective than prechlorination to reduce filter effluent turbidities. The cause of
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the improved coagulation is not clear, but several possibilities have been offered (Reckhow et al.,
1986), including:
• Oxidation of organics into more polar forms;
• Oxidation of metal ions to yield insoluble complexes such as ferric iron complexes; and
• Change in the structure and size of suspended particles.
2.2.7 Prevention of Algal Growth in Sedimentation Basins and
Filters
Prechlorination is often used to minimize operational problems associated with biological growth in
water treatment plants (AWWA, 1990; Culp/Wesner/Culp, 1986). Prechlorination will prevent
slime formation on filters, pipes, and tanks, and reduce potential taste and odor problems associated
with such slimes. Many sedimentation and filtration facilities operate with a small chlorine residual
to prevent growth of algae and bacteria in the launders and on the filter surfaces. This practice has
increased in recent years as utilities take advantage of additional contact time in the treatment units to
meet disinfection requirements under the SWTR.
2.2.8 Removal of Color
Free chlorine is used for color removal. A low pH is favored. Color is caused by humic compounds,
which have a high potential for DBF formation. The chlorine dosage and kinetics for color removal
are best determined through bench studies.
2.3 Disinfection Byproducts and Disinfection Residuals
2.3.1 Types of DBFs and Disinfection Residuals
Table 2-8 is a list, compiled by EPA, of DBFs and disinfection residuals that may be of health
concern. The table includes both the disinfectant residuals and the specific byproducts produced by
the disinfectants of interest in drinking water treatment. These contaminants of concern are grouped
into four distinct categories and include disinfectant residuals, inorganic byproducts, organic
oxidation byproducts, and halogenated organic byproducts. Tables 1-3 and 1-4 list the disinfection
byproducts and disinfectant residuals that are currently regulated.
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Table 2-8. List of Disinfection Byproducts and Disinfection Residuals
DISINFECTANT RESIDUALS
Free Chlorine
Hypocrjlorous Acid
Hypochlorite Ion
Chloramines
Monochloramine .
Chlorine Dioxide
INORGANIC BYPRODUCTS
Chlorate Ion
Chlorite Ion
Bromate Ion
lodate Ion
Hydrogen Peroxide
Ammonia
ORGANIC OXIDATION BYPRODUCTS
Aldehydes
Formaldehyde
Acetaldehyde
Glyoxal
Hexanal
Heptanal
Carboxylic Acids
Hexanoic Acid
Heptanoic Acid
Oxalic Acid
Assimilable Organic Carbon
HALOGENATED ORGANIC BYPRODUCTS
Trihalomethanes
Chloroform
Bromodichloromethane
Dibromochloromethane
Bromoform
Haloacetic Acids
Monochloroacetic Acid
Dichloroacetic Acid
Trichloroacetic Acid
Monobromoacetic Acid
Dibromoacetic Acid
Haloacetonitriles
Dichloroacetronitrile
Bromochloroacetonitrile
Dibromoacetonitrile
Trichloroacetonitrile
Haloketones
1,1-Dichloropropanone
1,1,1-Trichloropropanone
Chlorophenols
2-Chlorophenol
2,4-Dichlorophenol
2,4,6-Trichlorophenol
Chloropicrin
Chloral Hydrate
Cyanogen Chloride
N-Organochloramines
MX*
*3-Chloro-4-(dichloromethyl)-5-hydroxy-2(5H)-furanone
The production of DBFs depend on the type of disinfectant, the presence of organic material (e.g.,
TOC), bromide ion, and other environmental factors as discussed in this manual. By removing DBF
precursors, the formation of DBFs can be reduced.
The health effects of DBFs and disinfectants are generally evaluated with epidemiological studies
and/or toxicological studies using laboratory animals. Table 2-9 indicates the cancer classifications
of both disinfectants and DBFs as of January 1999. The classification scheme used by EPA is
shown at the bottom of Table 2-9. The EPA classification scheme for carcinogenicity weighs both
animal studies and epidemiologic studies, but places greater weight on evidence of carcinogenicity in
humans.
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B2
Table 2-9. Status of Health Information for Disinfectants and DBFs
Chloroform
Bromodichloromethane
Dibromochloromethane
Bromoform
Monochloroacetic Acid
Dichloroacetic Acid
Trichloroacetic Acid
Dichloroacetonitrile
Bromochloroacetonitrile
Dibromoacetonitrile
Trichloroacetonitrile
1,1 -Dichloropropanone
1,1,1-Trichloropropanone
2-Chlorophenol
2,4-Dichlorophenol
2,4,6-Trichlorophenol
Chloropicrin
Chloral Hydrate
Cyanogen Chloride
Formaldehyde
Chlorate
Chlorite
Bromate
Ammonia
Hypochlorous Acid
Hypochlorite
Monochloramine
Chlorine Dioxide
D
_,
B2
B1
(2)
D
B2
D
D
(1) The scheme for categorizing chemicals according to their carcinogenic potential is as follows:*
Group A:
Human Carcinogen
Group B;
Probable Human Carcinogen
Group C:
Possible Human Carcinogen
Group D:
Not Classifiable
Group E;
No Evidence of Carcinogenicity
(or Humans
Sufficient evidence in epidemiologic studies to support casual association
between exposure and cancer.
Limited evidence in epidemiologic studies (Group B1) and/or sufficient
evidence from animal studies (Group B2)
Limited evidence from animal studies and inadequate or no data in humans
Inadequate or no human and animal evidence of Carcinogenicity
No evidence of Carcinogenicity in at least two adequate animal tests in
different species or in adequate epidemiologic and animal studies.
* EPA is in tha process of revising the Cancer Guidelines
Source; USEPA. 1996
p> Based on Inhalation exposure.
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2. DISINFECTANT USE IN WA TER TREA TMENT
2.3.2 Disinfection Byproduct Formation
Halogenated organic byproducts are formed when natural organic matter (NOM) reacts with free
chlorine or free bromine. Free chlorine can be introduced to water directly as a primary or secondary
disinfectant, with chlorine dioxide, or with chloramines. Free bromine results from the oxidation of
the bromide ion in source water. Factors affecting formation of halogenated DBFs include type and
concentration of natural organic matter, oxidant type and dose, time, bromide ion concentration, pH,
organic nitrogen concentration, and temperature. Organic nitrogen significantly influences the
formation of nitrogen containing DBFs such as the haloacetonitriles, halopicrins, and cyanogen
halides (Reckhow et al., 1990; Hoigne and Bader, 1988). The parameter TOX represents the
concentration of total organic halides in a water sample (calculated as chloride). In general, less than
50 percent of the TOX content has been identified, despite evidence that several of these unknown
halogenated byproducts of water chlorination may be harmful to humans (Singer and Chang, 1989).
Non-halogenated DBFs are also formed when strong oxidants react with organic compounds found in
water. Ozone and peroxone oxidation of organics leads to the production of aldehydes, aldo- and
keto-acids, organic acids, and, when bromide ion is present, brominated organics (Singer, 1992).
Many of the oxidation byproducts are biodegradable and appear as biodegradable dissolved organic
carbon (BDOC) and assimilable organic carbon (AOC) in treated water.
Bromide ion plays a key role in DBF formation. Ozone or free chlorine oxidizes bromide ion to
hypobromate ion/hypobromous acid, which subsequently forms brominated DBFs. Brominated
organic byproducts include compounds such as bromoform, brominated acetic acids and
acetonitriles, bromopicrin, and cyanogen bromide. Only about one third of the bromide ions
incorporated into byproducts has been identified.
2.3.2.1 Disinfection Byproduct Precursors
Numerous researchers have documented that NOM is the principal precursor of organic DBF
formation (Stevens et al., 1976; Babcock and Singer 1979; Christman et al., 1983). Chlorine reacts
with NOM to produce a variety of DBFs, including THMs, haloacetic acids (HAAs), and others.
Ozone reacts with NOM to produce aldehydes, organic acids, and aldo- and keto-acids; many of
these are produced by chlorine as well (Singer and Harrington, 1993).
Natural waters contain mixtures of both humic and nonhumic organic substances. NOM can be
subdivided into a hydrophobic fraction composed of primarily humic material, and a hydrophilic
fraction composed of primarily fulvic material.
The type and concentration of NOM are often assessed using surrogate measures. Although
surrogate parameters have limitations, they are used because they may be measured more easily,
rapidly, and inexpensively than the parameter of interest, often allowing on-line monitoring of the
operation and performance of water treatment plants. Surrogates used to assess NOM include:
• Total and dissolved organic carbon (TOC and DOC);
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• Specific ultraviolet light absorbance (SUVA), which is the absorbance at 254 nm wavelength
(UV-254) divided by DOC (SUVA = (UV-254/DOC)*100, in L/mg-m);
• THM formation potential (THMFP) - a test measuring the quantity of THMs formed with a high
dosage of free chlorine and a long reaction time; and
• TTHM Simulated Distribution System (SDS) - a test to predict the TTHM concentration at some
selected point in a given distribution system, where the conditions of the chlorination test
simulate the distribution system at the point desired.
On average, about 90 percent of the TOC is dissolved. DOC is defined as the TOC able to pass
through a 0.45 urn filter. UV absorbance is a good technique for assessing the presence of DOC
because DOC primarily consists of humic substances, which contain aromatic structures that absorb
light in the UV spectrum. Oxidation of DOC reduces the UV absorbance of the water due to
oxidation of some of the organic bonds that absorb UV absorbance. Complete mineralization of
organic compounds to carbon dioxide usually does not occur under water treatment conditions;
therefore, the overall TOC concentration usually is constant.
DBF concentrations vary seasonally and are typically greatest in the summer and early fall for
several reasons:
* The rate of DBF formation increases with increasing temperature (Singer et al., 1992);
• The nature of organic DBF precursors varies with season (Singer et al., 1992); and
• Due to warmer temperatures, chlorine demand may be greater during summer months requiring
higher dosages to maintain disinfection.
If the bromide ion is present in source waters, it can be oxidized to hypobromous acid that can react
with NOM to form brominated DBFs, such as bromoform. Furthermore, under certain conditions,
ozone may react with the hypobromite ion (OBr") to form bromate ion (BrO"3).
The ratio of bromide ion to the chlorine dose affects THM formation and bromine substitution of
chlorine. Increasing the bromide ion to chlorine dose ratio shifts the speciation of THMs to produce
more brominated forms (Krasner et al., 1989; Black et al., 1996). In the Krasner et al. study, the
chlorine dose was roughly proportional to TOC concentration. As TOC was removed through the
treatment train, the chlorine dose decreased and TTHM formation declined. However, at the same
time, the bromide ion to chlorine dose increased, thereby shifting TTHM concentrations to the more
brominated THMs. Therefore, improving the removal of NOM prior to chlorination can shift the
speciation of halogenated byproducts toward more brominated forms.
Chloropicrin is produced by the chlorination of humic materials in the presence of nitrate ion
(Duguet et al., 1985; Thibaud et al., 1987). Thibaud et al. (1988) chlorinated humic compounds in
the presence of bromide ion to demonstrate the formation of brominated analogs to chloropicrin.
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2. DISINFECTANT USE IN WATER TREATMENT
2.3.2.2 Impacts ofpH on DBP Formation
The pH of water being chlorinated has an impact on the formation of halogenated byproducts as
shown in Table 2-10 (Reckhow and Singer, 1985; Stevens et al., 1989). THM formation increases
with increasing pH. Trichloroacetic acid, dichloroacetonitrile, and trichloropropanone formation
decrease with increased pH. Overall TOX formation decreases with increasing pH.
Based on chlorination studies of humic material in model systems, high pH tends to favor chloroform
formation over the formation of trichloroacetic acid and other organic halides. Accordingly, water
treatment plants practicing precipitative softening at pH values greater than 9.5 to 10 are likely to
have a higher fraction of TOX attributable to THMs than plants treating surface waters by .
conventional treatment in pH ranges of 6 to 8 (Singer and Chang, 1989).
Since the application of chlorine dioxide and chloramines may introduce free chlorine into water,
chlorination byproducts that may be formed would be influenced by pH as discussed above. Ozone
application to bromide ion containing waters at high pH favors the formation of bromate ion, while
application at low pH favors the formation of brominated organic byproducts. See discussion under
individual disinfectants for a more detailed discussion on pH impacts on DBP formation.
The pH also impacts enhanced coagulation (i.e., for ESWTR compliance) and Lead and Copper Rule
Compliance. These issues are addressed in EPA's Microbial and Disinfection Byproduct
Simultaneous Compliance Guidance Manual (expected to be available in 1999).
2.3.2.3 Organic Oxidation Byproducts
^-
Organic oxidation byproducts are formed by reactions between NOM and all oxidizing agents added
during drinking water treatment. Some of these byproducts are halogenated, as discussed in the
previous section, while others are not. The types and concentrations of organic oxidation byproducts
produced depend on the type and dosage of the oxidant being used, chemical characteristics and
concentration of the NOM being oxidized, and other factors such as the pH and temperature.
Specific chemical byproducts belonging to the classification of halogenated organic oxidation
products are listed in Table 2-10. As presented in Table 2-10, the formation of DBPs is pH
dependent. Comparisons in the table are made to the formation of TTHMs at a pH of 7.0. AOC is
not a specific organic contaminant, but a generally used surrogate measure of bacterial regrowth
potential in distribution systems. AOC is comprised of many chemical species, including the
aldehydes and carboxylic acids listed in Table 2-8. AOC formation studies, primarily performed in
the Netherlands, have shown that both ozonation and chlorination can increase concentrations of
AOC. This increase in AOC concentration is believed to be the result of oxidizing high molecular
weight organics to smaller and more readily bioassimilable molecules. Because AOC is not a
specific chemical contaminant, no specific health effects are attributable to AOC.
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Table 2-10. Conditions of Formation of DBFs
By-product
Chlorination at pH 5.0
Conditions of Formation
Chlorination at pH 7.0
Chlorination at pH 9.4
Total Trihalomethanes
Trichloroacetic Acid
Dichloroacetic Acid
Monochloroacetic Acid
Dibromoacetic Acid
Chloral Hydrate
Chloropicrin
Dichloroacetonitrile
Bromochloroacetonitrile
Dibromoacetonitrile
Trichloroacetonitrile
1,1,1-Trichloropropanone
Lower formation
Similar formation to that at
pH7.0
Similar formation to that at
pH 5.0 and 9.4 - perhaps
slightly higher at pH 7.0
Similar formation to that at
pHS.O
Similar formation to that
at pH 5.0 and 9.4-
perhaps slightly higher at
pH7.0
Higher formation
Lower formation
Similar formation to that at
pH 5.0 and 7.0 - perhaps
slightly higher at pH 7.0
At concentrations <5 ng/L,
trends not discernible
At concentrations <1 u.g/L,
trends not discernible
Similar formation to that at
pH7.0
At concentrations <1 u.g/L,
trends not discernible
At concentrations <5
trends not discernible
At concentrations <1 u.g/L,
trends not discernible
Similar formation to that at
pHS.O
At concentrations <5
trends not discernible
At concentrations <1
trends not discernible
Forms within 4 hours;
decays over time to <5
H9/L
At concentrations <1
trends not discernible
Higher formation
Forms within 4 hours;
decays over time to <5
At concentrations <1 ng/L,
trends not discernible
Concentrations <2 ng/L,
trends not discernible
At concentrations <2 u.g/L, At concentrations <2 u.g/L, At concentrations <2 u.g/L,
trends not discernible trends not discernible trends not discernible
At concentrations
<.5 ng/L, trends not
discernible
Not detected
Higher formation
At concentrations
<.5 ng/L, trends not
discernible
Not detected
At concentrations
<.5 u,g/L, trends not
discernible
Not detected
At concentrations <2 ng/L, Not detected
trends not discernible
Source: Stevens et al., 1989.
2.3.2.4 Inorganic Byproducts and Disinfectants
Table 2-11 shows some of the inorganic DBFs that are produced or remain as residual during
disinfection. As discussed earlier, bromide ion reacts with strong oxidants to form bromate ion and
other organic DBFs. Chlorine dioxide and chloramines leave residuals that are of concern for health
considerations, as well as for taste and odor. The significance of these compounds is discussed
further in subsequent chapters.
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2. DISINFECTANT USE IN WA TER TREA TMENT
Table 2-11. Inorganic DBFs Produced During Disinfection
Disinfectant Inorganic Byproduct or Disinfectant Residual Discussed
Chlorine Dioxide Chlorine Dioxide, Chlorite ion, Chlorate ion, Bromate ion (in presence of light)
Ozone Bromate ion, Hydrogen Peroxide
Chloramination Monochloratmine, Dichloramine, Trichloramine, Ammonia, Cyanogen Chloride
2.3.3 DBP Control Strategies
In 1983, the EPA identified technologies, treatment techniques, and plant modifications that
community water systems could use to comply with the maximum contaminant level for TTHMs.
The principal treatment modifications involved moving the point of chlorination downstream in the
water treatment plant, improving the coagulation process to enhance the removal of DBP precursors,
and using chloramines to supplement or replace the use of free chlorine (Singer, 1993). Moving the
point of chlorination downstream in the treatment train often is very effective in reducing DBP
formation, because it allows the NOM precursor concentration to be reduced during treatment prior
to chlorine addition. Replacing prechlorination by preoxidation with an alternate disinfectant that
produces less DBFs is another option for reducing formation of chlorinated byproducts.
Other options to control the formation of DBFs include; source water quality control, DBP precursor
removal, and disinfection strategy selection. An overview of each is provided below.
2.3.3.1 Source Water Quality Control
Source water control strategies involve managing the source water to lower the concentrations of
NOM and bromide ion in the source water. Research has shown that algal growth leads to the
production of DBP precursors (Oliver and Shindler, 1980; Wachter and Andelman, 1984; Karimi and
Singer, 1991). Therefore, nutrient and algal management is one method of controlling DBP
formation potential of source waters. Control of bromide ion in source waters may be accomplished
by preventing brine or salt water intrusion into the water source.
2.3.3.2 DBP Precursor Removal
Raw water can include DBP precursors in both dissolved and particulate forms. For the dissolved
precursors to be removed in conventional treatment, they must be converted to particulate form for
subsequent removal during settling and filtering. The THM formation potential generally decreases
by about 50 percent through conventional coagulation and settling, indicating the importance of
moving the point of chlorine application after coagulation and settling (and even filtration) to control
TOX as well as TTHM formation (Singer and Chang, 1989). Conventional systems can lower the
DBP formation potential of water prior to disinfection by further removing precursors with enhanced
coagulation, GAC adsorption, or membrane filtration prior to disinfection. Precursor removal
efficiencies are site-specific and vary with different source waters and treatment techniques.
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2. DISINFECTANT USE IN WATER TREATMENT
Aluminum (alum) and iron (ferric) salts can remove variable amounts of NOM. For alum, the
optimal pH for NOM removal is in the range of 5.5 to 6.0. The addition of alum decreases pH and
may allow the optimal pH range to be reached without acid addition. However, waters with very low
or very high alkalinities may require the addition of base or acid to reach the optimal NOM
coagulation pH (Singer, 1992).
GAC adsorption can be used following filtration to remove additional NOM. For most applications,
empty bed contact times in excess of 20 minutes are required, with regeneration frequencies on the
order of 2 to 3 months (Singer, 1992). These long contact times and frequent regeneration
requirements make GAC an expensive treatment option. In cases where prechlorination is practiced,
the chlorine rapidly degrades GAC. Addition of a disinfectant to the GAC bed can result in specific
reactions in which previously absorbed compounds leach into the treated water.
Membrane filtration has been shown effective in removing DBF precursors in some instances. In
pilot studies, ultrafiltration (UF) with a molecular weight cutoff (MWCO) of 100,000 daltons was
ineffective for controlling DBF formation. However, when little or no bromide ion was present in
source water, nanofiltration (NF) membranes with MWCOs of 400 to 800 daltons effectively
controlled DBF formation (Laine et al., 1993). In waters containing bromide ion, higher bromoform
concentrations were observed after chlorination of membrane permeate (compared with raw water).
This occurs as a result of filtration removing NOM while concentrating bromide ions in the permeate
thus providing a higher ratio of bromide ions to NOM than in raw water. This reduction in chlorine
demand increases the ratio of bromide to chlorine, resulting in higher bromoform concentrations after
chlorination of NF membrane permeate (compared with the raw water). TTHMs were lower in
chlorinated permeate than chlorinated raw water. However, due to the shift in speciation of THMs
toward more brominated forms, bromoform concentrations were actually greater in chlorinated
treated water than in chlorinated raw. water. Use of spiral-wound NF membranes (200-300 daltons)
more effectively controlled the formation of brominated THMs, but pretreatment of the water was
necessary (LSin6 et al., 1993). Significant limitations in the use of membranes are disposal of the
waste brine generated, fouling of membranes, cost of membrane replacement, and increasing energy
cost.
The promulgated DBPR requires enhanced coagulation as an initial step for removal of DBF
precursors. In addition to meeting MCLs and MRDLs, some water suppliers also must meet
treatment requirements to control the organic material (DBF precursors) in the raw water that
combines with disinfectant residuals to form DBFs. Systems using conventional treatment are
required to control precursors (measured as TOC) by using enhanced coagulation or enhanced
softening. A system must remove a specified percentage of TOC (based on raw water quality) prior
to the point of continuous disinfection (Table 2-12).
Systems using ozone followed by biologically active filtration or chlorine dioxide that meet specific
criteria would be required to meet the TOC removal requirements prior to addition of a residual
disinfectant. Systems able to reduce TOC by a specified percentage level have met the DBPR
treatment technique requirement.
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2. DISINFECTANT USE IN WA TER TREA TMENT
Table 2-12. Required Removal of TOC by Enhanced Coagulation for Surface
Water Systems* Using Conventional Treatment** (percent reduction)
Source Water
TOC (mg/L)
>2.0-4.0
>4.0-8.0
>8.0
0-60
35.0
45.0
50.0
Source Water Alkalinity (mg/L as CaCO3)
>60-120
25.0
35.0
40.0
>120~
15.0
25.0
30.0
+ Also applies to utilities that treat ground water under the influence of surface water.
++ Systems meeting at least one of the conditions in 40 CFR §§ 141.135(a)(1)(l)-(iv) are not required to operate with enhanced
coagulation.
+++ Systems practicing precipitative softening must meet the TOC removal requirements in this column.
If the system does not meet the percent reduction, it must determine its alternative minimum TOC
removal level. The primacy agency approves the alternative minimum TOC removal possible for the
system on the basis of the relationship between coagulant dose and TOC in the system based on
results of bench or pilot-scale testing. Enhanced coagulation is determined in part as the coagulant
dose where an incremental addition of 10 mg/L of alum (or an equivalent amount of ferric salt)
results in a TOC removal below 0.3 mg/L.
2.3.3.3 Disinfection Strategy Selection
In addition to improving the raw or predisinfectant water quality, alternative disinfection strategies
can be used to control DBFs. These strategies include the following:
• Use an alternative or supplemental disinfectant or oxidant such as chloramines or chlorine
dioxide that will produce fewer DBFs;
• Move the point of chlorination to reduce TTHM formation and, where necessary, substitute
chloramines, chlorine dioxide, or potassium permanganate for chlorine as a preoxidant;
• Use two different disinfectants or oxidants at various points in the treatment plant to avoid DBF
formation at locations where precursors are still present in high quantities;
• Use of powdered activated carbon for THM precursor or TTHM reduction seasonally or
intermittently; and
• Maximize precursor removal.
2.3.4 CT Factor
One of the most important factors for determining or predicting the germicidal efficiency of any
disinfectant is the CT factor, a version of the Chick-Watson law (Chick, 1908; Watson, 1908). The
CT factor is defined as the product of the residual disinfectant concentration, C, in mg/L, and the
/ contact time, T, in minutes, that residual disinfectant is in contact with the water.
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EPA developed CT values for the inactivation of Giardia and viruses under the SWTR. Table 2-13
compares the CT values for virus inactivation using chlorine, chlorine dioxide, ozone, chloramine,
and ultraviolet light disinfection under specified conditions. Table 2-14 shows the CT values for
inactivation of Giardia cyst using chlorine, chloramine, chlorine dioxide, and ozone under specified
conditions. The CT values shown in Table 2-13 and Table 2-14 are based on water temperatures of
10°C and pH values in the range of 6 to 9. CT values for chlorine disinfection are based on a free
chlorine residual. Note that chlorine is less effective as pH increases from 6 to 9. In addition, for a
given CT value, a low C and a high T is more effective than the reverse (i.e., a high C and a low T).
For all disinfectants, as temperature increases, effectiveness increases.
Table 2-13. CT Values for Inactivation of Viruses
Disinfectant
Chlorine1
Chloramine2
Chlorine Dioxide3
Ozone
UV
Units
mg • min/L
mg • min/L
mg • min/L
mg • min/L
mW • s/cm2
2-log
3
643
4.2
0.5
21
Inactivation
3-log
4
1,067
12.8
0.8
36
4-log
6
1,491
25.1
1.0
not available
CT values were obtained from AWWA, 1991.
1 Values are based on a temperature of 10°C, pH range of 6 to 9, and a free chlorine residual of 0.2 to 0.5 mg/L.
1 Values are based on a temperature of 10°C and a pH of 8.
3 Values are based on a temperature of 10°C and a pH range of 6 to 9.
Table 2-14. CT Values for Inactivation of Giardia Cysts
Disinfectant
Chlorine'
Chloramine2
Chlorine Dioxide3
Ozone3
0.5-log
17
310
4
0.23
1-log
35
615
7.7
0.48
Inactivation <
1.5-log
52
930
12
0.72 -
(mg • min/L)
2-log
69
1,230
15
0.95
2.5-log
87
1,540
19
1.2
3-log
104
1,850
23
1.43
CT values were obtained from AWWA, 1991.
1 Values are based on a free chlorine residual less than or equal to 0.4 mg/L, temperature of 10°C, and a pH of 7.
* Values are based on a temperature of 10°C and a pH in the range of 6 to 9.
3 Values are based on a temperature of 10°C and a pH of 6 to 9.
2.4 Pathogen Inactivation Versus DBF Formation
Table 2-15 presents a summary of disinfection parameter impacts on pathogen inactivation and DBF
formation.
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2. DISINFECTANT USE IN WATER TREATMENT
Table 2-15. Summary of Disinfection Impacts
Disinfection Parameter
Disinfectant Type
Disinfectant Strength
Disinfectant Dose
Type of Organism
Contact Time
PH
Temperature
Turbidity
Dissolved Organics
Typical Impact on Pathogen
Inactivation
Typical Impact on DBP Formation
Depends on inactivation efficacy
The stronger the disinfectant, the
quicker the disinfection process.
Increasing the disinfectant dose
increases the disinfection rate.
Susceptibility to disinfection varies
according to pathogen group: In general,
protozoa are more resistant to
disinfectants than bacteria and viruses.
Depends on disinfectant reactivity
The stronger the disinfectant, the
greater the amount of DBFs.
Increasing the disinfectant dose typically
increases the rate of DBP formation.
None.
Increasing the contact time decreases
the disinfectant dose required for a
given level of inactivation.
Increasing contact time with an
equivalent disinfectant dose increases
the formation of DBPs.
pH may affect the disinfectant form and,
in-turn, the efficiency of the disinfectant.
The impact of pH varies with DBP. See
Section 2.3.2.3 for a brief summary of
relationships between pH and DBP
formation.
Increasing the temperature increases
the rate of disinfection.
Increasing temperature is typically
associated with faster oxidation kinetics,
hence, increased DBP formation.
Particles responsible for turbidity can
surround and shield pathogenic
microorganisms from disinfectants.
Dissolved organics can interfere with
disinfection by creating a demand and
reducing the amount of disinfectant
available for pathogen inactivation.
Increased turbidity may be associated
with increased NOM, which represents
an increased amount of DBP precursors
for the formation of DBPs when
disinfectant is applied.
Increased dissolved organics will
represent a larger amount of DBP
precursor for the formation of DBPs
when disinfectant is applied.
2.5 Disinfectant Residual Regulatory Requirements
One of the most important factors for evaluating the merits of alternative disinfectants is their ability
to maintain the microbial quality in the water distribution system. Disinfectant residuals may serve
to protect the distribution system against regrowth (Snead et al., 1980). The SWTR requires that
filtration and disinfection must be provided to ensure that the total treatment of the system achieves
at least a 3-log removal/inactivation of Giardia cysts and a 4-log removal/inactivation of viruses. In
addition, the disinfection process must demonstrate by continuous monitoring and recording that the
disinfectant residual in the water entering the distribution system is never less than 0.2 mg/L for more
than 4 hours.
Several of the alternative disinfectants examined in this manual cannot be used to meet the residual
requirements stated in the SWTR. For example, if either ozone or ultraviolet light disinfection are
used as the primary disinfectant, a secondary disinfectant such as chlorine or chloramines should be
utilized to obtain a residual in the distribution system.
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DBF formation continues in the distribution system due to reactions between the residual disinfectant
and organics in the water. Koch et al. (1991) found that with a chlorine dose of 3-4 mg/L, THM and
HAA concentrations increase rapidly during the first 24 hours in the distribution system. After the
initial 48 hours, the subsequent increase in THMs is very small. Chloral hydrate concentrations
continued to increase after the initial 24 hours, but at a reduced rate. Haloketones actually decreased
in the distribution system.
Nieminski et al. (1993) evaluated DBF formation in the simulated distribution systems of treatment
plants in Utah. Finished water chlorine residuals ranged from 0.4 to 2.8 mg/L. Generally, THM
values in the distribution system studies increased by 50 to 100 percent (range of 30 to 200 percent)
of the plant effluent value after 24-hour contact time. The 24-hour THM concentration was
essentially the same as the 7-day THM formation potential. HAA concentrations in the simulated
distribution system was about 100 percent (range of 30 to 200 percent) of the HAA in the plant
effluent. The 7-day HAA formation potential was sometimes higher, or below the distribution
system values. If chlorine is used as a secondary disinfectant, one should therefore anticipate a 100-
percent increase in the plant effluent THMs, or plan to reach the 7-day THM formation level in the
distribution system.
2.6 Summary of Current National Disinfection Practices
Most water treatment plants disinfect water prior to distribution. The 1995 Community Water
Systems Survey (USEPA, 1997a) reports that 81 percent of all community water systems provide
some form of treatment on all or a portion of their water sources. The survey also found that
virtually all surface water systems provide some treatment of their water. Of those systems reporting
no treatment, 80 percent rely on ground water as their only water source.
The most commonly used disinfectants/oxidants are chlorine, chlorine dioxide, chloramines, ozone,
and potassium permanganate. Table 2-16 shows a breakdown on the chemical usage from the
survey. Note that the table shows the percentages of systems using the particular chemical as either
disinfectant or some other role. The table shows the predominance of chlorine in surface and ground
water disinfection treatment systems with more than 60 percent of the treatment systems using
chlorine as disinfectant/oxidant. Potassium permanganate on the other hand, is used by many
systems, but its application is primarily for oxidation, rather than for disinfection.
Permanganate will have some beneficial impact on disinfection since it is a strong oxidant that will
reduce the chemical demand for the ultimate disinfection chemical. Chloramine is used by some
systems and is more frequently used as a post-treatment disinfectant.
The International Ozone Association conducted a survey of ozone facilities in the United States
(IOA, 1997). The survey documented the types of ozone facilities, size, objective of ozone
application, and year of operation. Table 2-17 summarizes the findings from the survey. The most
common use for ozone is for oxidation of iron and manganese, and for taste and odor control.
Twenty-four of the 158 ozone facilities used GAC following ozonation. In addition to the 158
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2. DISINFECTANT USE IN WA TER TREA TMENT
operating ozone facilities, the survey identified 19 facilities under construction and another 30 under
design. The capacity of the systems range from less than 25 gpm to exceeding 500 mgd. Nearly half
of the operating facilities have a capacity exceeding 1 mgd. Rice et al. (1998) found that as of May
1998, 264 drinking water plants in the United States are using ozone.
Table 2-16. Disinfection Practices of Water Systems that Include
Some Form of Treatment
Service Population
Treatment
<100
101-500
501-
1,000
1,001-
3,300
3,301-
10,000
10,001-
50,000
50,001-
100,000
Over
100,001
Total
Surface Water Systems
Total Number of
Systems
218
432
330
845
679
626
103
104
3,337
Pre-Disinfection, Oxidation/Softening
Chlorine
Chlorine dioxide
Chloramines
Ozone
KMnCX,
Predisinfection/oxidation
Lime/Soda ash softening
Recarbonation
Post-Disinfection
Chlorine
Chlorine dioxide
Chloramines
Postdisinfection
combinations
59.0%
0
4.6
0
0
0
6.8
0
49.7 %
0
0
0
73.9%
0
0
0
4.9
0
9.8
0
51.6%
0
0
0
67.3%
0
1.1
0
9.6 .
2.0
20.9
0
80.6%
0
0
0
66.3%
5.0
2.1
0
9.9
2.9
16.2
0
62.8%
0
2.9
2.1
68.8%
4.7
0
0.3
15.2
0.6
14.3
2.1
77.9%
0.3
2.1
4.0
58.6%
13.2
2.2
0
28.3
9.2
11.7
4.7
71.1%
4.9
15.6
3.9
47.5%
14.2
15.5
5.4
25.9
5.1
3.5
0.6
73.8 %
5.9
29.4
1.9
57.1%
7.8
10.8
5.8
28.5
4.3
5.9
6.3
63.6 %
11.2
24.2
1.6
63.8%
6.3
3.1
0.9
16.0
3.5
12.5
1.9
67.5 %
1.6
8.1
3.0
Ground Water Systems
Total Number of
Systems
9,042
10,367
4,443
4,422
2,035
1,094
120
56
31,579
Pre-Oisinfection, Oxidation/Softening
Chlorine
Chlorine dioxide
Chloramines
Ozone
KMnO4
Predisinfection/oxidation
Lime/Soda ash softening
Recarbonation
Post-Disinfection
Chlorine
Chlorine dioxide
Chloramines
Postdisinfection
combinations
64.2 %
1.3
0
0
0
0.3
2.9
0
23.0 %
0
0
0
69.9 %
0
0
0
0.9
0.5
2.9
0.5
23.4%
1.0
0
0
56.7 %
0
0
0
2.2
0
2.2
0
32.5%
' 0
0
0
73.2 %
0
0
0
0.6
0.7
3.6
0.6
28.3 %
0
0
0
60.6 %
0
0
0
5.8
1.0
3.5
1.4
42.5%
0
0.1
0.1
57.4%
0
0.6
0
'3.2
2.6
3.8
1.5
41.9%
0.6
1.1
0.1
36.2 %
3.1
1.4
0
7.0
0
5.0
2.8
54.5 %
0
3.9
0
38.1 %
0
0.7
0.6
0
0
9.1
1.1
65.8 %
0
4.3
0
63.9 %
0.3
0.1
0
1.8
0.7
3.2
0.6
31.0%
0.4
0.3
0
Source: USEPA, 1997a.
April 1999
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Table 2-17. Ozone Application in Water Treatment Plants in the United States
Ozone Objective
THM Control
Disinfection
Iron/Manganese, Taste and Odor Control
Total
Number of Plants
50
63
92
158
% Plants
32
40
58
-
Source: IOA, 1997.
2.7 Chlorine
Although chlorine is not a focus of this guidance manual, the following section provides a brief
overview of chlorine use in the water treatment industry to compare with the alternative disinfectants
discussed in this manual. Since there is a wealth of excellent literature on chlorine's uses and
performance capabilities, summarizing this large body of knowledge here is neither practical nor
necessary (see, for example: White, 1992; Chlorine Institute, 1996; and Connell, 1996).
One of the recent developments in chlorine disinfection is the use of multiple and interactive
disinfectants. In these applications, chlorine is combined with a second disinfectant to achieve
improved disinfection efficiency and/or effective DBF control. A detailed discussion on multiple
disinfectants, including chlorine combinations, is provided in Chapter 9.
As described earlier, the 1995 Community Water System Survey (USEPA, 1997a), indicated that the
majority of all surface water and ground water systems in the United States use chlorine for
disinfection.
Chlorine has many attractive features that contribute to its wide use in the industry. Four of the key
attributes of chlorine are that it:
* Effectively inactivates a wide range of pathogens commonly found in water;
• Leaves a residual in the water that is easily measured' and controlled;
• Is economical; and
• Has an extensive track record of successful use in improving water treatment operations (despite
the dangers associated with chlorine application and handling, specifically chlorine gas, it still
maintains an excellent safety record).
There are, however, some concerns regarding chlorine usage that may impact its uses such as:
• Chlorine reacts with many naturally occurring organic and inorganic compounds in water to
produce undesirable DBFs;
• Hazards associated with using chlorine, specifically chlorine gas, require special treatment and
response programs; and
* High chlorine doses can cause taste and odor problems.
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Chlorination is used in water treatment facilities primarily for disinfection. Because of chlorine's
oxidizing powers, it has been found to serve other useful purposes in water treatment, such as (White,
1992):
• Taste and odor control;
• Prevention of algal growths;
• Maintenance of clear filter media;
• Removal of iron and manganese;
• Destruction of hydrogen sulfide;
• Bleaching of certain organic colors;
• Maintenance of distribution system water quality by controlling slime growth;
• Restoration and preservation of pipeline capacity;
• Restoration of well capacity, water main sterilization; and
• Improved coagulation by activated silica.
2.7.1 Chlorine Chemistry
Chlorine for disinfection typically is used in one of three forms: chlorine gas, sodium hypochlorite,
or calcium hypochlorite. A brief description of the chemistry of these three chemicals is provided
below.
2.7.1.1 Chlorine Gas
Chlorine gas hydrolyzes rapidly in water to form hypochlorous acid (HOCl). The following equation
presents the hydrolysis reaction:
C/2(g) + H2O => HOCl + H+ Cr Equation 1
Note that the addition of chlorine gas to water reduces the pH of the water due to the production of
hydrogen ion.
Hypochlorous acid is a weak acid (pKa of about 7.5), meaning it dissociates slightly into hydrogen
and hypochlorite ions as noted in Equation 2:
HOCl <=> H+ + O Of Equation 2
Between a pH of 6.5 and 8.5 this dissociation is incomplete and both HOCl and OC1" species are
present to some extent (White, 1992). Below a pH of 6.5, no dissociation of HOCl occurs, while
above a pH of 8.5, complete dissociation to OC1" occurs. As the germicidal effects of HOCl is much
higher than that of OC1", chlorination at a lower pH is preferred.
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2.7.1.2 Hypochlorite
In addition to chlorine gas, chlorine is also available in hypochlorite form as both aqueous solutions
and dry solids. The most common aqueous hypochlorite solution is sodium hypochlorite. The most
common form of dry solid hypochlorite is calcium hypochlorite (White, 1992).
Sodium Hypochlorite. Sodium hypochlorite is produced when chlorine gas is dissolved in a sodium
hydroxide solution. Sodium hypochlorite solution typically contains 12.5 percent available chlorine
(White, 1992). One gallon of 12.5 percent sodium hypochlorite solution typically contains the
equivalent of one pound of chlorine.
The reaction between sodium hypochlorite and water is shown in the following reaction:
NaOCl + H2O => HOC/ + Na+ + OH' Equations
Equation 3 shows that the application of sodium hypochlorite to water produces hypochlorous acid,
similar to chlorine gas hydrolysis (Equation 1). However, unlike chlorine hydrolysis, the addition of
sodium hypochlorite to water yields a hydroxyl ion that will increase the pH of the water. In
addition, excess sodium hydroxide is used to manufacture sodium hypochlorite, which will further
increase the pH of the water.
Calcium Hypochlorite. Calcium hypochlorite is formed from the precipitate that results from
dissolving chlorine gas in a solution of calcium oxide (lime) and sodium hydroxide. Granular
calcium hypochlorite commercially available typically contains 65 percent available chlorine. This
means that 1.5 pounds of calcium hypochlorite contains the equivalent of one pound of chlorine. The
reaction between calcium hypochlorite and water is shown in the following reaction:
Ca(OCl)2+2E2O=> 2HOC/ + Ca+++2OH- Equation 4
Equation 4 shows that the application of calcium hypochlorite to water also produces hypochlorous
acid, similar to chlorine gas hydrolysis (Equation 1). Similar to sodium hypochlorite solution, the
addition of calcium hypochlorite to water yields hydroxyl ions that will increase the pH of the water.
2.7.2 Chlorine Generation
Onsite generation of chlorine has recently become practical. These generation systems, using only
salt and electric power, can be designed to meet disinfection and residual standards and to operate
unattended at remote sites. Considerations for chlorine generation include cost, concentration of the
brine produced, availability of raw materials, and the reliability of the process (AWWA and ASCE,
1997).
2.7.2.1 Chlorine
Chlorine gas can be generated by a number of processes including the electrolysis of alkaline brine or
hydrochloric acid, the reaction between sodium chloride and nitric acid, or the oxidation of
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hydrochloric acid. About 70 percent of the chlorine produced in the United States is manufactured
from the electrolysis of salt brine and caustic solutions in a diaphragm cell (White, 1992). Since
chlorine is a stable compound, it is typically produced off-site by a chemical manufacturer. Once
produced, chlorine is packaged as a liquefied gas under pressure for delivery to the site in railcars,
tanker trucks, or cylinders.
2.7.2.2 Sodium Hypochlorite
Dilute sodium hypochlorite solutions (less than 1 percent) can be generated electrochemically on-site
from salt brine solution. Typically, sodium hypochlorite solutions are referred to as liquid bleach or
Javelle water. Generally, the commercial or industrial grade solutions produced have hypochlorite
strengths of 10 to 16 percent. The stability of sodium hypochlorite solution depends on the
hypochlorite concentration, the storage temperature, the length of storage (time), the impurities of the
solution, and exposure to light. Decomposition of hypochlorite over time can affect the feed rate and
dosage, as well as produce undesirable byproducts such as chlorite ions or chlorate (Gordon et al.,
1995). Because of the storage problems, many systems are investigating onsite generation of
hypochlorite in lieu of its purchase from a manufacturer or vendor (USEPA, 1998b).
2.7.2.3 Calcium Hypochlorite
To produce calcium hypochlorite, hypochlorous acid is made by adding chlorine monoxide to water
and then neutralizing it with a lime slurry to create a solution of calcium hypochlorite. The water is
removed from the solution, leaving granulated calcium hypochlorite. Generally, the final product
contains up to 70 percent available chlorine and 4 to 6 percent lime. Storage of calcium hypochlorite
is a major safety consideration. It should never be stored where it is subject to heat or allowed to
contact any organic material of an easily oxidized nature (USEPA, 1998b).
2.7.3 Primary Uses and Points of Application of Chlorine
2.7.3.1 Uses
The main usage of chlorine in drinking water treatment is for disinfection. However, chlorine has
also found application for a variety of other water treatment objectives such as, the control of
nuisance organisms, oxidation of taste and odor compounds, oxidation of iron and manganese, color
removal, and as a general treatment aid to filtration and sedimentation processes (White, 1992;
Connell, 1996; Culp/Wesner/Culp, 1986). Table 2-18 presents a summary of chlorine uses and
doses.
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Table 2-18. Chlorine Uses and Doses
Application
Iron
Manganese
Biological
growth
Taste/odor
Color removal
Zebra mussels
Asiatic clams
Typical Dose
0.62 mg/mg Fe
0.77 mg/mg Mn
1-2 mg/L
Varies
Varies
2-5 mg/L
0.2-0.5 mg/L(a)
0.3-0.5 mg/L(a)
Optimal pH
7.0
7-8
9.5
6-8
6-8
4.0-6.8
Reaction Time
less than 1 hour
1-3 hour
minutes
NA
Varies
Minutes
Shock level
Maintenance level
Continuous
Effectiveness
Good
Slow kinetics
Good
Varies
Good
Good
Good
Other
Considerations
Reaction time
increases at
lower pH
DBP formation
Effectiveness
depends on
compound
DBP formation
DBP formation
DBP formation
Notes:
(i| Residual, not dose
Sources! Adapted in part Irom White, 1992; Connell, 1996; Cutp/Wesner/Culp, 1986.
2.7.3.2 Points of Application
At conventional surface water treatment plants, chlorine is typically added for prechlorination at
either the raw water intake or flash mixer, for intermediate chlorination ahead of the filters, for
postchlorination at the filter clearwell, or for rechlorination of the distribution system (Connell,
1996). Table 2-19 summarizes the typical uses for each point of application.
Table 2-19. Typical Chlorine Points of Application and Uses
Point of Application
Typical Uses
Raw Water Intake
Flash Mixer (prior to sedimentation)
Filter Influent
Filter Clearwell
Distribution System
Zebra mussel and Asiatic clam control, control biological growth
Disinfection, iron and manganese oxidation, taste and odor control,
oxidation of hydrogen sulfide
Disinfection, control biological growth in filter, iron and manganese
oxidation, taste and odor control, algae control, color removal
Disinfection
Maintain disinfectant residual
Sources: Connell, 1996; White, 1992; AWWA, 1990.
2.7.3.3 Typical Doses
Table 2-20 shows the typical dosages for the various forms of chlorine. The wide range of chlorine
gas dosages most likely represents its use as both an oxidant and a disinfectant. While sodium
hypochlorite and calcium hypochlorite can also serve as both an oxidant and a disinfectant, their
higher cost may limit their use.
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2. DISINFECTANT USE IN WA TER TREA TMENT
Table 2-20. Typical Chlorine Dosages at Water Treatment Plants
Chlorine Compound Range of Doses
Calcium hypochlorite 0.5-45 mg/L
Sodium hypochlorite 0.2-2 mg/L
Chlorine gas 1-16 mg/L
3, as adapted from EPA's review of public water systems' Initial Sampling Plans which were
required by EPA's Information Collection Rule (ICR)
2.7.4 Pathogen Inactivation and Disinfection Efficacy
2.7.4.1 Inactivation Mechanisms
Research has shown that chlorine is capable of producing lethal events at or near the cell membrane
as well as affecting DNA. In bacteria, chlorine was found to adversely affect cell respiration,
transport, and possibly DNA activity (Haas and Engelbrecht, 1980). Chlorination was found to cause
an immediate decrease in oxygen utilization in both Escherichia coli and Candida parapsilosis
studies. The results also found that chlorine damages the cell wall membrane, promotes leakage
through the cell membrane, and produces lower levels of DNA synthesis for Escherichia coli,
Candida parapsilosis, and Mycobacterium fortuitum bacteria. This study also showed that chlorine
inactivation is rapid and does not require bacteria reproduction (Haas and Engelbrecht, 1980). These
observations rule out mutation or lesions as the principal inactivation mechanisms since these
mechanisms require at least one generation of replication for inactivation to occur.
2.7.4.2 Environmental Effects
Several environmental factors influence the inactivation efficiency of chlorine, including water
temperature, pH, contact time, mixing, turbidity, interfering substances, and the concentration of
available chlorine. In general, the highest levels of pathogen inactivation are achieved with high
chlorine residuals, long contact times, high water temperature, and good mixing, combined with a
low pH, low turbidity, and the absence of interfering substances. Of the environmental factors, pH
and temperature have the most impact on pathogen inactivation by chlorine. The effect of pH and
temperature on pathogen inactivation are discussed below.
pH. The germicidal efficiency of hypochlorous acid (HOC1) is much higher than that of the
hypochlorite ion (OC1~). The distribution of chlorine species between HOC1 and OC1" is determined
by pH, as discussed above. Because HOC1 dominates at low pH, chlorination provides more
effective disinfection at low pH. At high pH, OC1" dominates, which causes a decrease in
disinfection efficiency.
The inactivation efficiency of gaseous chlorine and hypochlorite is the same at the same pH after
chlorine addition. Note, however, that addition of gaseous chlorine will decrease the pH (see
Equation 1) while the addition of hypochlorite will increase the pH of the water (see Equation 3 and
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Equation 4). Therefore, without pH adjustment to maintain the same treated water pH, gaseous
chlorine will have greater disinfection efficiency than hypochlorite.
The impact of pH on chlorine disinfection has been demonstrated in the field. For example, virus
inactivation studies have shown that 50 percent more contact time is required at pH 7.0 than at pH
6.0 to achieve comparable levels of inactivation. These studies also demonstrated that a rise in pH
from 7.0 to 8.8 or 9.0 requires six times the contact time to achieve the same level of virus
inactivation (Gulp and Gulp, 1974). Although these studies found a decrease in inactivation with
increasing pH, some studies have shown the opposite effect. A 1972 study reported that viruses were
more sensitive to free chlorine at high pH than at low pH (Scarpino et al., 1972).
Temperature. For typical drinking water treatment temperatures, pathogen inactivation increases
with temperature. Virus studies indicate that the contact time should be increased by two to three
times to achieve comparable inactivation levels when the water temperature is lowered by 10°C
(Clarke et al., 1962).
2.7.4.3 Disinfection Efficacy
Since its introduction, numerous investigations have been made to determine the germicidal
efficiency of chlorine. Although there are widespread differences in the susceptibility of various
pathogens, the general order of increasing chlorine disinfection difficulty are bacteria, viruses, and
then protozoa.
Bacteria Inactivation. Chlorine is an extremely effective disinfectant for inactivating bacteria. A
study conducted during the 1940s investigated the inactivation levels as a function of time for E. coli,
Pseudomonas aeruginosa, Salmonella typhi, and Shigella dysenteriae (Butterfield et al., 1943).
Study results indicated that HOC1 is more effective than OCl'for inactivation of these bacteria.
These results have been confirmed by several researchers that concluded that HOC1 is 70 to 80 times
more effective than OCl"for inactivating bacteria. (Culp/Wesner/Culp, 1986).
Virus Inactivation. Chlorine has been shown to be a highly effective viricide. One of the most
comprehensive virus studies was conducted in 1971 using treated Potomac estuary water (Liu et al.,
1971). The tests were performed to determine the resistance of 20 different enteric viruses to free
chlorine under constant conditions of 0.5 mg/L free chlorine and a pH and temperature of 7.8 and
2°C, respectively. In this study, the least resistant virus was found to be reovirus and required 2.7
minutes to achieve 99.99 percent inactivation (4 log removal). The most resistant virus was found to
be a poliovirus, which required more than 60 minutes for 99.99 inactivation. The corresponding CT
range required to achieve 99.99 percent inactivation for all 20 viruses was between 1.4 to over 30
mg-min/L.
Virus survival studies have also been conducted on a variety of both laboratory and field strains
(AWWA, 1979). All of the virus inactivation tests in this study were performed at a free chlorine
residual of 0.4 mg/L, a pH of 7.0, a temperature of 5°C, and contact times of either 10, 100, or 1,000
minutes. Test results showed that of the twenty cultures tested only two poliovirus strains reached
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2. DISINFECTANT USE IN WA TER TREA TMENT
99.99 percent inactivation after 10 minutes (CT = 4 mg-min/L), six poliovirus strains reached 99.99
percent inactivation after 100 minutes (CT = 40 mg-min/L), and 11 of the 12 polioviruses plus one
Coxsackievirus strain (12 out of a total of 20 viruses) reached 99.99 percent inactivation after 1,000
minutes (CT = 400 mg-min/L).
Protozoa Inactivation. Chlorine has been shown to have limited success inactivating protozoa.
Data obtained during a 1984 study indicated that the resistance of Giardia cysts are two orders of
magnitude higher than that of enteroviruses and more than three orders of magnitude higher than the
enteric bacteria (Hoff et al., 1984). CT requirements for Giardia cysts inactivation when using
chlorine as a disinfectant has been determined for various pH and temperature conditions (AWWA,
1991). These CT values increase at low temperatures and high pH (See also Table 2-13).
Chlorine has little impact on the viability of Cryptosporidium oocysts when used at the relatively low
doses encountered in water treatment (e.g., 5 mg/L). Approximately 40 percent removals (0.2 log) of
Cryptosporidium were achieved at CT values of both 30 and 3,600 mg-min/L (Finch et al., 1994).
Another study determined that "no practical inactivation was observed" when oocysts were exposed
to free chlorine concentrations ranging from 5 to 80 mg/L at pH 8, a temperature of 22°C, and
contact times of 48 to 245 minutes (Gyiirek et al., 1996). CT values ranging from 3,000 to 4,000
mg-min/L were required to achieve 1-log of Cryptosporidium inactivation at pH 6.0 and temperature
of 22°C. During this study, one trial in which oocysts were exposed to 80 mg/L of free chlorine for
120 minutes was found to produce greater than 3-logs of inactivation.
2.7.4.4 CT Curves
Chlorine is regarded as a strong disinfectant that is effective at inactivating bacteria and viruses, and
under certain circumstances, Giardia. Because of chlorine's extremely high virus inactivation
efficiency, CT values are almost always governed by protozoa inactivation. For example, Figure 2-1
shows the CT values required to achieve between 0.5 and 3-logs of virus and Giardia inactivation
(AWWA, 1991). As shown, the CT values required to achieve the recommended disinfection
efficiency for conventional filtration systems (i.e., 0.5-log Giardia cyst and 2-log virus inactivation
level) are 23 and 3 mg min/L, respectively.
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Log Pathogen Niaclivatlon Ltvtl
Figure 2-1. Free Chlorine Giardia and Virus CT Requirements
CT values for Giardia inactivation for various pH values and temperatures at a chlorine dose of
3.0 mg/L are shown in Figures 2-2 and 2-3. As shown, the inactivation efficacy of free chlorine
decreases with increasing pH and/or decreasing temperature. CT values shown in Figures 2-2 and 2-
3 are based on animal infectivity and excystation studies. CT values ranging from 0.5 to 3-log
inactivation at temperatures of 0.5 and 5°C were based on a multiplicative model, and applying first
order kinetics to the 99 percent upper confidence interval of the 99.99-percentile CT values. CT
values for temperatures above 5°C were estimated by assuming a twofold decrease for every 10°C
decrease in temperature.
pH(au)
Figure 2-2. CT Values for Inactivation of Giardia Cysts by Free Chlorine at 10°C
(at CI2 dose of 3.0 mg/L)
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15
Temperature (°C)
Figure 2-3. CT Values for Inactivation of Giardia Cysts by Free Chlorine at pH 7.0
(at CI2 dose of 3.0 mg/L)
2.7.5 DBP Formation and Control
2.7.5.1 DBP Formation
Halogenated organics are formed when natural organic matter (NOM) reacts with free chlorine or
free bromine. Free chlorine is normally introduced into water directly as a primary or secondary
disinfectant. Free bromine results from the oxidation by chlorine of the bromide ion in the source
water. Factors affecting the formation of these halogenated DBFs include type and concentration of
NOM, chlorine form and dose, time, bromide ion concentration, pH, organic nitrogen concentration,
and temperature. Organic nitrogen significantly influenced the formation of nitrogen containing
DBFs, including haloacetonitriles, halopicrines, and cyanogen halides (Reckhow et al., 1990; Hoigne
and Bader, 1988).
The formation of DBFs is strongly related to TOC at the point of disinfection. DBP formation also
correlates with the amount of chlorine consumed (Singer et al., 1995). Stevens et al. (1989) found
that higher TTHM formation occurs at high pH (9.4) than at low pH (5.0) while HAA showed no
clear trend as a function of pH. A survey of 35 water utilities conducted by MWDSC (Krasner et al.,
1989) showed the median TTHM and HAA concentrations measured as 39 and 19 ug/L (i.e. more
THMs than HAAs are formed). However, a subsequent study by Singer et al. (1995) found a
reversal in dominance, with more HAAs than THMs produced in waters from North Carolina
utilities. They postulated that the reason for this change is due to the lower pH levels and differences
in TOC and bromide concentrations in the North Carolina waters. Pourmoghaddas et al. (1993)
showed that brominated and mixed brominated/chlorinated THMs and HAAs are formed when using
chlorine in the presence of bromide.
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The occurrence of THMs and HAAs is important because regulatory limits are placed on both groups
of compounds. One water utility may therefore find that its chlorination practice is limited by the
production of THMs, while another will find that HAAs limit the use of chlorine. This distribution
of THMs and HAA is a function of the TOC and bromide concentration in the water, as well as the
pH during chlorination.
Of note, Chlorate is produced as a byproduct when hypochlorite degrades during storage.
2.7.5.2 DBF Control
DBFs can be controlled by several means, including removing the DBF precursors, modifying the
chlorination strategy, changing disinfectants, or removing the DBF itself. Because DBFs are
difficult to remove once they are formed, control strategies typically focus on the first three
methods.
Studies have shown that removal of TTHM precursors tends to remove the formation potential for
the other DBFs. Generally, aggregate DBF formation will decrease as the removal of TOC increases.
Recent research indicates that moving the point of chlorination back into the treatment process can
reduce the formation of DBFs.
Summers et al. (1997) recently summarized the results from four studies evaluating the impact of
pretreatment on DBF formation. Jar tests were conducted to simulate the water treatment through
rapid mix, coagulation, flocculation, and sedimentation. Chlorine was added at various points in the
jar testing to simulate the impact of various dose points on production of DBFs. The results clearly
demonstrate the benefits of delaying the point of chlorination downstream in the treatment train to
take advantage of precursor removal during initial flocculation and sedimentation processes. Table
2-21 summarizes the results from this study.
Table 2-21. Percent Reduction in DBP Formation by Moving
Chiorination Point Later In Treatment Train
Chlorination point
Pre rapid mix
Post rapid mix
Mid flocculation
Post sedimentation
TTHM
Baseline (%)
Basis
1.6
8.7
21
TTHM
Enhanced (%)
17
21
36
48
HAAS
Baseline (%)
Basis
5.3
14
35
HAAS
Enhanced (%)
4.7
21
36
61
Notes: Source: USEPA, 1997b based on Summers et al., 1997.
Baseline » Baseline coagulant (alum) dose for optimal turbidity removal (-30 mg/L)
Enhanced = Enhanced coagulant (alum) dose for optimal TOC removal (-52 mg/L)
Table 2-21 also shows the benefit of enhanced coagulation to reduce DBP production. The THM
reduction of 21 percent by moving the chlorination point to post sedimentation is more than doubled
to 48 percent by enhanced coagulation. The HAA removal increases from 45 to 61 percent under
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enhanced coagulation with post sedimentation chlorination. Therefore, DBF control by selecting the
optimal dose location and conditions along with enhanced precursor removal can significantly reduce
DBF formation at low added cost.
White (1992) suggested that pretreatment goals should include: 1) maximizing THM precursor
removal; 2) reducing ammonia-N concentration to 0.10 mg/L; 3) reducing organic-N concentration to
0.05 mg/L; and 4) limiting 15-minute chlorine demand to 0.5 mg/L. These guidelines should
improve raw water quality sufficiently to allow the use of the free chlorine residual process without
exceeding the EPA MCLs for TTHMs.
2.7.6 Operational Considerations
2.7.6.1 Application Methods
Different application methods are used, depending upon the form of chlorine used. The following
paragraphs describe the typical application methods for chlorine, sodium hypochlorite, and calcium
hypochlorite.
Chlorine. Liquefied chlorine gas is typically evaporated to gaseous chlorine prior to metering. The
heat required for evaporation can be provided through either a liquid chlorine evaporator or the
ambient heat input to the storage container. Once the compressed liquid chlorine is evaporated,
chlorine gas is typically fed under vacuum conditions. Either an injector or a vacuum induction
mixer usually creates the required vacuum. The injector uses water flowing through a venturi to
draw the chlorine gas into a side stream of carrier water to form a concentrated chlorine solution.
This solution is then introduced into the process water through a diffuser or mixed with a mechanical
mixer. A vacuum induction mixer uses the motive forces of the mixer to create a vacuum and draws
the chlorine gas directly into the process water at the mixer.
Sodium Hypochlorite. Sodium hypochlorite solutions degrade over time. For example, a 12.5
percent hypochlorite solution will degrade to 10 percent in 30 days under "best case" conditions
(White, 1992). Increased temperature, exposure to light, and contact with metals increase the rate of
sodium hypochlorite degradation (Connell, 1996).
Sodium hypochlorite solution is typically fed directly into the process water using a type of metering
pump. Similar to chlorine solution, sodium hypochlorite is mixed with the process water with either
a mechanical mixer or induction mixer. Sodium hypochlorite solution is typically not diluted prior to
mixing to reduce scaling problems.
Calcium Hypochlorite. Commercial high-level calcium hypochlorite contains at least 70%
available chlorine (USEPA, 1991). Under normal storage conditions, calcium hypochlorite loses 3 to
5% of its available chlorine in a year (AWWA and ASCE, 1997). Calcium hypochlorite comes in
powder, granular, and compressed tablet forms (USEPA, 1991). Typically, calcium hypochlorite
solution is prepared by mixing powdered or granular calcium hypochlorite with a small flow. The
highly chlorinated solution is then flow paced into drinking water flow.
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2.7.6.2 Safety and Handling Considerations
Chlorine. Chlorine gas is a strong oxidizer. The U.S. Department of Transportation classifies
chlorine as a poisonous gas (Connell, 1996). Fire codes typically regulate the storage and use of
chlorine. In addition, facilities storing more than 2,500 pounds of'chlorine are subject to the
following two safety programs:
• Process Safety Management standards regulated by the Occupational Safety and Health
Administration under 29 CFR 1910.
• The Risk Management Program Rule administered by EPA under Section 112(r) of the Clean Air
Act.
All of these regulations (as well as local and state codes and regulations) must be considered during
the design and operation of chlorination facilities at a water treatment plant.
Sodium Hypochlorite. Sodium hypochlorite solution is a corrosive liquid with an approximate pH
of 12 (AWWA, 1990). Therefore, typical precautions for handling corrosive materials such as
avoiding contact with metals, including stainless steel, should be used.
Sodium hypochlorite solutions may contain chlorate. Chlorate is formed during the both the
manufacturing and storage of sodium hypochlorite (due to degradation of the product). Chlorate
formation can be minimized by reducing the degradation of sodium hypochlorite (Gilbert et al.,
1995) by limiting storage time, avoid high temperatures and reduce light exposure.
Spill containment must be provided for the sodium hypochlorite storage tanks. Typical spill
containment structures include containment for the entire contents of the largest tank (plus freeboard
for rainfall or fire sprinklers), no uncontrolled floor drains, and separate containment areas for each
incompatible chemical.
Calcium Hypochlorite. Calcium hypochlorite is an oxidant and as such should be stored separately
from organic materials that can be readily oxidized. It should also be stored away from sources of
heat. Improperly stored calcium hypochlorite has caused spontaneous combustion fires (White,
1992).
2.8 Summary
2.8.1 Advantages and Disadvantages of Chlorine Use
The following list presents selected advantages and disadvantages of using chlorine as a disinfection
method for drinking water (Masschelein, 1992; Process Applications, Inc., 1992). Because of the
wide variation of system size, water quality, and dosages applied, some of these advantages and
disadvantages may not apply to a particular system.
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Advantages
• Oxidizes soluble iron, manganese, and sulfides
• Enhances color removal
• Enhances taste and odor
• May enhance coagulation and filtration of particulate contaminants
• Is an effective biocide
• Is the easiest and least expensive disinfection method, regardless of system size
• Is the most widely used disinfection method, and therefore, the best known
• Is available as calcium and sodium hypochlorite. Use of these solutions is more advantageous
for smaller systems than chlorine gas because they are easier to use, are safer, and need less
equipment compared to chlorine gas
• Provides a residual.
Disadvantages
• May cause a deterioration in coagulation/filtration of dissolved organic substances
• Forms halogen-substituted byproducts
• Finished water could have taste and odor problems, depending on the water quality and dosage
• Chlorine gas is a hazardous corrosive gas
• Special leak containment and scrubber facilities could be required for chlorine gas
• Typically, sodium and calcium hypochlorite are more expensive than chlorine gas
• Sodium hypochlorite degrades over time and with exposure to light
• Sodium hypochlorite is a corrosive chemical
• Calcium hypochlorite must be stored in a cool, dry place because of its reaction with moisture
and heat
• A precipitate may form in a calcium hypochlorite solution because of impurities, therefore, an
antiscalant chemical may be needed
• Higher concentrations of hypochlorite solutions are unstable and will produce chlorate as a
byproduct
• Is less effective at high pH
• Forms oxygenated byproducts that are biodegradable and which can enhance subsequent
biological growth if a chlorine residual is not maintained.
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Release of constituents bound in the distribution system (e.g., arsenic) by changing the redox
state.
2.8.2 Summary Table
Table 2-22 presents a summary of the considerations for the use of chlorine as a disinfectant.
Table 2-22. Summary of Chlorine Disinfection
Description
Chlorination may be performed using chlorine gas or other chlorinated
compounds that may be in liquid or solid form. Chlorine gas can be
generated by a number of processes including the electrolysis of alkaline
brine or hydrochloric acid, the reaction between sodium chloride and nitric
acid, or the oxidation of hydrochloric acid. Since chlorine is a stable
compound, chlorine gas, sodium hypochlorite, and calcium hypochlorite are
typically produced off-site by a chemical manufacturer.
Consideration
Generation
Primary uses
Inactivation efficiency
Byproduct formation
Point of application
Special considerations
The primary use of Chlorination is disinfection. Chlorine also serves as an
oxidizing agent for taste and odor control, prevention of algal growths,
maintaining clear filter media, removal of iron and manganese, destruction
of hydrogen sulfide, color removal, maintaining the water quality at the
distribution systems, and improving coagulation.
The general order of increasing chlorine disinfection difficulty is bacteria,
viruses, and then protozoa. Chlorine is an extremely effective disinfectant
for inactivating bacteria and highly effective viricide. However, chlorine is
less effective against Giardia cysts. Cryptosporidium oocysts are highly
resistant to chlorine.
When added to the water, free chlorine reacts with NOM and bromide to
form DBPs, primarily THMs, some haloacetic acids (HAAs), and others.
Raw water storage, precoagulation/post-raw water storage,
presedimentation/ postcoagulation, postsedimentation/prefiltration, post
filtration (disinfection), or in the distribution system.
Because chlorine is such a strong oxidant and extremely corrosive, special
storage and handling considerations should be considered in the planning
of a water treatment plant. Additionally, health concerns associated with
handling and use of chlorine is an important consideration.
2.8.3 Reference for Additional Information on Chlorine
With the focus of this manual on disinfectants other than chlorine, all of chlorine's uses and
capabilities are not described here. For more detailed information regarding the use of chlorine in
water treatment, refer to the list of references provided below. For complete references, see the
References section at the end of this chapter.
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• AWWA(1990)
• Connell (1996)
• DeMers and Renner (1992)
• Hazen and Sawyer (1992)
• Hoigne and Bader( 1988)
• Sawyer etal. (1994)
• Singer (1988)
• White (1992)
2.9 References
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3. AWWA (American Water Works Association). 1991. Guidance Manual for Compliance with
the Filtration and Disinfection Requirements for Public Works Systems using Surface Water
Sources.
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Identification and Treatment. AWWA, Denver, CO.
5. AWWA and ASCE (American Water Works Association and American Society of Civil
Engineers). 1997. Water Treatment Plant Design. McGraw-Hill, New York, NY.
6. AWWARF (American Water Works Association Research Foundation) and Lyonnaise des
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7. Babcock, D.S. and P.C. Singer. 1979. "Chlorination and Coagulation of Humic and Fulvic
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8. Beneson, A.S. 1981. "Control of Communicable Diseases in Man." APHA.
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10. Black, B.D.,, G.W. Harrington, and P.C. Singer. 1996. "Reducing Cancer Risks by Improving
Organic Carbon Removal." J. AWWA. 88(6):40.
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11. Brady, Thomas J., I.E. Van Benschoten, and J.N. Jensen. 1996. "Chlorination Effectiveness for
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26. Culp/Wesner/Culp. 1986. Handbook of Public Water Systems. Van Nostrand Reinhold, New
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30. Farvardin, M.R. and A.G. Collins. 1990. Mechanism(s) of Ozone-Induced Coagulation of
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31. Ferguson, D.W., J.T. Gramith, and MJ. McGuire. 1991. "Applying Ozone for Organics Control
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38. Haas C.N. and R.S. Engelbrecht. 1980. "Physiological Alterations of Vegetative
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46. Karimi, A.A. and P.C. Singer. 1991. Trihalomethane.Formation in Open Reservoirs. /. AWWA.
83 (3):84.
47. Klerks, P.L. and P.C. Fraleigh, P.C. 1991. "Controlling Adult Zebra Mussels with Oxidants."
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48. Koch, B., S.W. Krasner, MJ. Sclimenti, and W.K. Schimpff. 1991. "Predicting the Formation of
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51. Laine, J.M., J.G. Jacangelo, E.W. Cummings, K.E. Cams, J. Mallevialle. 1993. "Influence of
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64. Pourmoghaddas, H., A.A. Stevens, R.N. Kinman, R.C. Dressman, L.A. Moore, J.C. Ireland.
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78. Singer, P.C. 1988. Alternative Oxidant and Disinfectant Treatment Strategies for Controlling
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79. Singer P.C. 1992. "Formation and Characterization of Disinfection Byproducts." Presented at
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80. Singer P.C. 1993. "Trihalomethanes and Other Byproducts Formed From the Chlorination of
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82. Singer P.C., and G.W. Harrington. 1993. "Coagulation of DBP Precursors: Theoretical and
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84. Snead, M.C., et al. 1980. Benefits of Maintaining a Chlorine Residual in Water Supply Systems.
EPA 600/2-80-010. . - -
85. Stevens, A.A., et al. 1976. "Chlorination of Organics in Drinking Water." J. AWWA. 8(11):615.
86. Stevens, A.A., L.A. Moore, R.J. Miltner. 1989. "Formation and Control of Non-Trihalomethane
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87. Suffet, I. H., C. Anselme, and J. Mallevialle. 1986. "Removal of Tastes and Odors by
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91. Thibaud, H., J. DeLaat, and M. Dore". 1988. "Effects of Bromide Concentration on the
Production of Chloropicrin During Chlorination of Surface Waters: Formation of Brominated
Trihalonitromethanes." Water Res. 22(3):381.
92. Tobiason, J.E., J.K. Edzwald, O.D. Schneider, M.B. Fox, and H.J. Dunn. 1992. "Pilot Study of
the Effects of Ozone and PEROXONE on In-Line Direct Filtration." J. AWWA. 84(12):72-84.
93. USEPA. 1998a. Occurrence Assessment for Disinfectants and Disinfection Byproducts in Public
Drinking Water Supplies. Science Applications International Corporation under contract for
Office of Ground Water and Drinking Water. Washington, DC.
94. USEPA. 1998b. Technologies and Costs for Control of Disinfection Byproducts. Prepared by
Malcolm Pirnie, Inc for U.S. Environmental Protection Agency, Office of Ground Water and
Drinking Water, PB93-162998.
95. USEPA. 1997a. Community Water System Survey - Volumes I and II; Overview. EPA 815-R-97-
00la, -001 b. January.
96. USEPA. 1997b. "National Primary Drinking Water Regulations: Disinfectants and Disinfection
Byproducts; Notice of Data Availability; Proposed Rule." Federal Register. 62(212):59387-
59484. November 3.
97. USEPA. 1996. Drinking Water Regulations and Health Advisories. EPA 822-B-96-002,
October.
98. USEPA. 1991. Manual of Individual and Non-Public Works Supply Systems. Office of Water,
EPA 570/9-91-004.
99. Van Benschoten, I.E., J.N. Jensen, D. Harrington, and D.J. DeGirolamo. 1995. "Zebra Mussel
Mortality With Chlorine." J. AWWA. 87(5): 101-108.
100. Wachter, J.K., and J.B. Andelman. 1984. "Organohalide Formation on Chlorination of Algal
Extracellular Products." Env.Sci. Technol. 18(111):811.
101. Watson, H.E. 1908. "A Note on the Variation of the Rate of Disinfection With Change in the
Concentration of the Disinfectant." J. Hygiene.. 8:538.
102. White, G.C. 1992. Handbook of Chlorination and Alternative Disinfectants. Van Nostrand
Reinhold, New York, NY.
103. Witherell, L.E, R.W. Duncan, K.M. Stone, L.J. Stratton, L. Orciari, S. Kappel, D.A. Jillson.
1988. "Investigation of Legionella Pneumophila in Drinking Water." /. AWWA. 80 (2):88-93.
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3. OZONE
Ozone was first used for drinking water treatment in 1893 in the Netherlands. While being used
frequently in Europe for drinking water disinfection and oxidation, it was slow to transfer to the
United States. In 1987, the Los Angeles Aqueduct Filtration Plant was placed in service and now
treats up to 600 mgd of drinking water. In 1991, approximately 40 water treatment plants each
serving more than 10,000 people in the United States utilized ozone (Langlais et al., 1991). This
number has grown significantly, with Rice (in press) reporting that as of April 1998, 264 operating
plants in the United States use ozone. Most of these facilities are small: 149 plants are below 1 mgd.
Ozone is used in water treatment for disinfection and oxidation. Early application of ozone in the
United States was primarily for non-disinfection purposes such as color removal or taste and odor
control. However, since the implementation of the SWTR and proposal of the DBF rule, ozone
usage for primary disinfection has increased in the United States.
3.1 Ozone Chemistry
Ozone exists as a gas at room temperature. The gas is colorless with a pungent odor readily
detectable at concentrations as low as 0.02 to 0.05 ppm (by volume), which is below concentrations
of health concern. Ozone gas is highly corrosive and toxic.
Ozone is a powerful oxidant, second only to the hydroxyl free radical, among chemicals typically
used in water treatment. Therefore, it is capable of oxidizing many organic and inorganic
compounds in water. These reactions with organic and inorganic compounds cause an ozone demand
in the water treated, which should be satisfied during water ozonation prior to developing a
measurable residual.
Ozone is sparingly soluble in water. At 20°C, the solubility of 100 percent ozone is only 570 mg/L
(Kinman, 1975). While ozone is more soluble than oxygen, chlorine is 12 times more soluble than
ozone. Ozone concentrations used in water treatment are typically below 14 percent, which limits
the mass transfer driving force of gaseous ozone into the water. Consequently, typical concentrations
of ozone found during water treatment range from <0.1 to Img/L, although higher concentrations can
be attained under optimum conditions.
Basic chemistry research (Hoigne and Bader, 1983a and 1983b; Glaze et al., 1987) has shown that
ozone decomposes spontaneously during water treatment by a complex mechanism that involves the
generation of hydroxyl free radicals. The hydroxyl free radicals are among the most reactive
oxidizing agents in water, with reaction rates on the order of 1010 - 1013 M"1 s"1, approaching the
diffusion control rates for solutes such as aromatic hydrocarbons, unsaturated compounds, aliphatic
alcohols, and formic acid (Hoigne and Bader, 1976). On the other hand, the half-life of hydroxyl free
radicals is on the order of microseconds, therefore concentrations of hydroxyl free radicals can never
reach levels above 10~12 M (Glaze and Kang, 1988).
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3. OZONE .
As shown in Figure 3-1 ozone can react by either or both modes in aqueous solution (Hoigne and
Bader, 1977):
• Direct oxidation of compounds by molecular ozone (O3(aq)).
• Oxidation of compounds by hydroxyl free radicals produced during the decomposition of
ozone.
Direct Oxidation of Substrate
Ozone Decomposition
via *OH
Indirect Oxidation of Substrate
by Hydroxyl Radical
Radical Consumption by
HCO3-, CO3-2, etc.
• Byproducts
• Byproducts
Figure 3-1. Oxidation Reactions of Compounds (Substrate)
During Ozonation of Water
The two oxidation pathways compete for substrate (i.e., compounds to oxidize). The direct oxidation
with aqueous ozone is relatively slow (compared to hydroxyl free radical oxidation) but the
concentration of aqueous ozone is relatively high. On the other hand, the hydroxyl radical reaction is
fast, but the concentration of hydroxyl radicals under normal ozonation conditions is relatively small.
Hoigne" and Bader (1977) found that:
• Under acidic conditions, the direct oxidation with molecular ozone is of primary importance;
and
• Under conditions favoring hydroxyl free radical production, such as high pH, exposure to
UV, or addition of hydrogen peroxide, the hydroxyl oxidation starts to dominate.
This latter mechanism is used in advanced oxidation processes such as discussed in Chapter 7,
Peroxone, to increase the oxidation rates of substrates.
The spontaneous decomposition of ozone occurs through a series of steps. The exact mechanism and
reactions associated have not been established, but mechanistic models have been proposed (Hoigne
and Bader, 1983a and 1983b; Glaze, 1987). It is believed that hydroxyl radicals forms as one of the
intermediate products, and can react directly with compounds in the water. The decomposition of
ozone in pure water proceeds with hydroxyl free radicals produced as an intermediate product of
ozone decomposition, resulting in the net production of 1.5 mole hydroxyl free radicals per mole
ozone.
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3. OZONE
In the presence of many compounds commonly encountered in water treatment, ozone decomposition
forms hydroxyl free radicals. Ozone demands are associated with the following:
• Reactions with natural organic matter (NOM) in the water. The oxidation of NOM leads to
the formation of aldehydes, organic acids, and aldo- and ketoacids (Singer, 1992).
• Organic oxidation byproducts. Organic oxidation byproducts are generally more amenable to
biological degradation and can be measured as assimilable organic carbon (AOC) or
biodegradable dissolved organic carbon (BDOC).
• Synthetic organic compounds (SOCs). Some SOCs can be oxidized and mineralized under
favorable conditions. To achieve total mineralization, hydroxyl radical oxidation should
usually be the dominant pathway, such as achieved in advanced oxidation processes.
• Oxidation of bromide ion. Oxidation of bromide ion leads to the formation of hypobromous
acid, hypobromite ion, bromate ion, brominated organics, and bromamines (see Figure 3-2).
• Bicarbonate or carbonate ions, commonly measured as alkalinity, will scavenge the hydroxyl
radicals and form carbonate radicals (Staehelin et al., 1984; Glaze and Kang, 1988). These
reactions are of importance for advanced oxidation processes where the radical oxidation
pathway is predominant.
NH2Br
NH,
HOBr
POM fc Brominated
organic comp.
Br03-
BrO-
Br
Source: Gunten and Hoigne, 1996.
Figure 3-2. Reaction of Ozone and Bromide Ion Can Produce Bromate Ion and
Brominated Organics
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3.2 Ozone Generation
3.2.1 Ozone Production
Because ozone is an unstable molecule, it should be generated at the point of application for use in
water treatment. It is generally formed by, combining an oxygen atom with an oxygen molecule
(02):
3020203
This reaction is endothermic and requires a considerable input of energy.
Schonbein (Langlais et < biblio >) first discovered synthetic ozone through the electrolysis of
sulfuric acid. Ozone can be produced several ways, although one method, corona discharge,
predominates in the ozone generation industry. Ozone can also be produced by irradiating an
oxygen-containing gas with ultraviolet light, electrolytic reaction and other emerging technologies as
described by Rice (1996).
Corona discharge, also known as silent electrical discharge, consists of passing an oxygen-containing
gas through two electrodes separated by a dielectric and a discharge gap. Voltage is applied to the
electrodes, causing an electron flowthrough across the discharge gap. These electrons provide the
energy to disassociate the oxygen molecules, leading to the formation of ozone. Figure 3-3 shows a
basic ozone generator.
HEAT
HIGH VOLTAGE
ELECTRODE
AC
02
CORONA
DISCHARGE GAP
GROUND
ELECTRODE
HEAT
Figure 3-3. Basic Ozone Generator
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3.2.2 System Components
As shown in Figure 3-4, ozone water treatment systems have four basic components: a gas feed
system, an ozone generator, an ozone contactor, and an off-gas destruction system. The gas feed
system provides a clean, dry source of oxygen to the generator. The ozone contactor transfers the
ozone-rich gas into the water to be treated, and provides contact time for disinfection (or other
reactions). The final process step, off-gas destruction, is required as ozone is toxic in the
concentrations present in the off-gas. Some plants include an off-gas recycle system that returns the
ozone-rich off-gas to the first contact chamber to reduce the ozone demand in the subsequent
chambers. Some systems also include a quench chamber to remove ozone residual in solution.
Influent
o o o
o o
o o
„ o
o o o
o o
o o
Aozpne Gas
Atmosphere
Off-Gas
Ozone Destruction
Ozone
Quench
(Option)
Effluent
Figure 3-4. Simplified Ozone System Schematic
3.2.2.1 Gas Feed Systems
Ozone feed systems are classified as using air, high purity oxygen or mixture of the two. High purity
oxygen can be purchased and stored as a liquid (LOX), or it can be generated on-site through either a
•cryogenic process, with vacuum swing adsorption (VSA), or with pressure swing adsorption (PSA).
Cryogenic generation of oxygen is a complicated process and is feasible only in large systems.
Pressure swing adsorption is a process whereby a special molecular sieve is used under pressure to
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selectively remove nitrogen, carbon dioxide, water vapor, and hydrocarbons from air, producing an
oxygen rich (80-95 percent Oa) feed gas. The components used in pressure swing adsorption
systems are similar to high pressure air feed systems in that both use pressure swing molecular
absorption equipment. Low pressure air feed systems use a heat reactivated desiccant dryer.
Oxygen Feed Systems - Liquid oxygen feed systems are relatively simple, consisting of a storage
tank or tanks, evaporators to convert the liquid to a gas, filters to remove impurities, and pressure
regulators to limit the gas pressure to the ozone generators.
Air Feed Systems - Air feed systems for ozone generators are fairly complicated as the air should be
properly conditioned to prevent damage to the generator. Air should be clean and dry, with a
maximum dew point of -60° C (-80° F) and free of contaminants. Air preparation systems typically
consist of air compressors, filters, dryers, and pressure regulators. Figure 3-5 is a schematic of large
scale air preparation system.
Particles greater than 1 |J.m and oil droplets greater than 0.05 Jim should be removed by filtration
(Langlais et al., 1991). If hydrocarbons are present in the feed gas, granular activated carbon filters
should follow the particulate and oil filters. Moisture removal can be achieved by either compression
or cooling (for large-scale system), which lowers the holding capacity of the air, and by desiccant
drying, which strips the moisture from the air with a special medium. Desiccant dryers are required
for all air preparation systems. Large or small particles and moisture cause arcing which damages
generator dielectrics.
Typically, desiccant dryers are supplied with dual towers to allow regeneration of the saturated tower
while the other is in service. Moisture is removed from the dryer by either an external heat source or
by passing a fraction (10 to 30 percent) of the dried air through the saturated tower at reduced
pressure. Formerly, small systems that require only intermittent use of ozone , a single desiccant
tower is sufficient, provided that it is sized for regeneration during ozone decomposition time.
Air preparation systems can be classified by the operating pressure: ambient, low (less than 30 psig),
medium, and high (greater than 60 psig) pressure. The distinguishing feature between low and high
pressure systems is that high pressure systems can use a heatless dryer. A heatless dryer operates
normally in the 100 psig range, rather than the 60 psig range. Rotary lobe, centrifugal, rotary screw,
liquid ring, vane, and reciprocating compressors can be used in air preparation systems. Table 3-1
lists the characteristics of many of these types of compressors.
Reciprocating and liquid ring compressors are the most common type used in the United States,
particularly in small systems, the former because the technology is so prevalent and the latter because
liquid ring compressors do not need aftercoolers. Air receivers are commonly used to provide
variable air flow from constant volume compressors. Oil-less compressors are used in modern
systems to avoid hydrocarbons in the feed gas (Dimitriou, 1990).
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3.
Desiccant
Filter Compressor Aftercooler Separator Dryer
Air
Receiver
Dry, Clean Air
to Generator
Figure 3-5. Schematic of an Air Preparation System
Table 3-1. Types of Compressors Used in Air Preparation Systems
Compressor Type Pressure
Rotary lobe Low -15 psi
Centrifugal
Volume
Constant or variable with
unloading
Comments
Common in Europe
30 -100 psi depending Variable, high volume
on no. of stages
Medium efficiency, cost effective in high
volumes
Rotary Screw 50 psi (single stage) to Variable with unloading
100 psi (2 stage)
Liquid Ring
Vane
10-80 psi
High-to 100 psi
Constant volume
Constant or variable
Slightly more efficient than rotary lobe, draws
approximately 40% of full load power in
unloaded state, available in non-lubricated
desjgn for larger capacities.
Does not require lubrication or aftercooler,
relatively inefficient, common in United
States.
Relatively inefficient, not common in U.S.
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Table 3-2 presents a comparison of the advantages and disadvantages of each gas feed system.
Table 3-2. Comparison of Air and High Purity Oxygen Feed Systems
Source
Air
Oxygen (general)
LOX
Cryogenic Oxygen
Generation
Advantages
Commonly used equipment
Proven technology
Suitable for small and large systems
Higher ozone concentration (8-14%)
Approximately doubles ozone concentration
for same generator
Suitable for small and large systems
Less equipment required
Simple to operate and maintain
Suitable for small and intermediate systems
Can store excess oxygen to meet peak
demands
Equipment similar to air preparation systems
Feasible for large systems
Can store excess oxygen to meet peak
demands
Disadvantages
More energy consumed per ozone volume
produced
Extensive gas handling equipment required
Maximum ozone concentration of 3-5%
Safety concerns
Oxygen resistant materials required
• Variable LOX costs
• Storage of oxygen onsite (Fire Codes, i.e.
safety concerns)
• Loss of LOX in storage when not in use
• More complex than LOX
• Extensive gas handling equipment required
• Capital intensive
• Complex systems to operate and maintain
3.2.2.2 Ozone Generators
The voltage required to produce ozone by corona discharge is proportional to the pressure of the
source gas in the generator and the width of the discharge gap. Theoretically, the highest yield
(ozone produced per unit area of dielectric) would result from a high voltage, a high frequency, a
large dielectric constant, and a thin dielectric. However, there are practical limitations to these
parameters. As the voltage increases, the electrodes and dielectric materials are more subject to
failure. Operating at higher frequencies produces higher concentrations of ozone and more heat
requiring increased cooling to prevent ozone decomposition. Thin dielectrics are more susceptible to
puncturing during maintenance. The design of any commercial generator requires a balance of ozone
yield with operational reliability and reduced maintenance.
Two different geometric configurations for the electrodes are used in commercial ozone generators:
concentric cylinders and parallel plates. The parallel plate configuration is commonly used in small
generators and can be air cooled. Figure 3-6 shows the basic arrangement for the cylindrical
configuration. The glass dielectric/high voltage electrode in commercial generators resembles a
fluorescent light bulb and is commonly referred to as a "generator tube."
Most of the electrical energy input to an ozone generator (about 85 percent) is lost as heat (Rice,
1996). Because of the adverse impact of temperature on the production of ozone, adequate cooling
should be provided to maintain generator efficiency. Excess heat is removed usually by water
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3. OZONE
flowing around the stainless steel ground electrodes. The tubes are arranged in either a horizontal or
vertical configuration in a stainless steel shell, with cooling water circulating through the shell.
02
(AIR)
HIGH
DrVTTMTIAI /
r(J 1 C.N 1 IKL £.
ELECTRODE^
GAS
/-STAINLESS STEEL GROUND ELECTRODE
ooooooooooooooooooooooooooooooooooo
ooooooooooooooooooooooooooooooooooo
-»_
)"
m ,
ooooooooooooooooooooooooooooooooooo
ooooooooooooooooooooooooooooooooooo
— Oa
7 GLASS
DIELECTRIC
MATERIAL
^ A-
^
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3. OZONE
Table 3-3. Comparison of Primary Characteristics of Low, Medium, and High
Frequency Ozone Generators
Bubble Diffuser Contactors
The bubble diffuser contactor is commonly used for ozone contacting in the United States and
throughout the world (Langlais et al., 1991). This method offers the advantages of no additional
energy requirements, high ozone transfer rates, process flexibility, operational simplicity, and no
moving parts. Figure 3-7 illustrates a typical three stage ozone bubble diffuser contactor. This
illustration shows a countercurrent flow configuration (ozone and water flowing in opposite
directions), an alternating cocurrent/countercurrent arrangement, and a cocurrent flow configuration
(ozone and water flowing in the same direction). Also, the number of stages can vary from two to six
for ozone disinfection, with the majority of plants using two or three chambers for contacting and
reaction (Langlais et al., 1991).
Bubble diffuser contactors are typically constructed with 18 to 22 ft water depths to achieve 85 to 95
percent ozone transfer efficiency. Since all the ozone is not transferred into the water, the contactor
chambers are covered to contain the off-gas. Off-gas is routed to an ozone destruct unit, usually
catalysts, thermal, or thermal/catalysts.
Characteristic
Degree of Electronics Sophistication
Peak Voltages
Turndown Ratio
Cooling Water Required (gal/lb of
ozone produced)
Typical Application Range
Operating Concentrations
wt-%inair
wt - % in oxygen
Optimum Ozone Production (as a
proportion of total generator capacity)
Optimum Cooling Water Differential
Power Required (kW-h/Ib Oa)
Air Feed System Power Requirements
(kW-h/lb Oj)
Low Frequency
(50 -60 Hz)
low
~~"l9.5
__ ._._ ..
0.5 to 1.0
< 500 Ib/day
0.5 to 1.5%
2.0 to 5.0%
60 to 75%
8°to10°F
air feed: 8 to 12
Oz feed: 4 to 6
5to7
Medium Frequency
(up to 1,000 Hz)
high
11.5
10:1
0.5 to 1.5
to 2,000 !b/day
1.0to2.5%+
2 to 12%
90 to 95%
5° to 8°F
air feed: 8 to 12
Oa feed: 4 to 6
5 to 7
High Frequency
(> 1,000 Hz)
high
10
10:1
0.25 to 1
to 2,000 Ib/day
1.0 to 2.5%*
2 to 12%
90 to 95%
5° to 8°F
air feed: 8 to 12
Oz feed: 4 to 6
5 to 7
Source; Adapted from Rice, 1996, with modifications.
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3. OZONE
A. Counter Current Contactor
Influent
9
•4-
•I
> Ozone Gas
B. Counter and Cocurrent Contactor
Influent
.
• "
•/4'.
• « O
O 0
O f ^1
e '
•;r.
e o e
^*r* ^
\
v_
f
^
Ozone Gas
C. Cocurrent Contactor
Influent
•1
t
Ozone Gas -
Figure 3-7. Ozone Bubble Contactor
Bubble diffuser contactors use ceramic or stainless steel diffusers that are either rod-type or disc-type
to generate bubbles. Design considerations for these diffusers (Renner et al., 1988) include:
• Gas flow range of 0.5 to 4.0 scfm;
• Maximum headloss of 0.5 psig;
• Permeability of 2 to 15 cfm/ft2/in of diffuser thickness; and porosity of 35 to 45 percent.
The configuration of the bubble diffuser contactor structure should best be designed to provide plug
flow hydraulics. This configuration will minimize the overall volume of the contactor while still
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3. OZONE
meeting the CT requirements for the system. Contactor volume is determined in conjunction with the
applied ozone dosage and estimated residual ozone concentration to satisfy the disinfection CT
requirement.
Table 3-4 summarizes the advantages and disadvantages of the bubble diffuser contactor (Langlais et
al., 1991). Also, diffuser pore clogging can be a problem when ozone dosages are intermittent and/ or
when iron and manganese oxidation is required. Channeling of bubbles is dependent on the type of
diffusers used and the spacing between diffusers.
Table 3-4. Bubble Diffuser Contactor Advantages and Disadvantages
Advantages
Disadvantages
No moving parts
Effective ozone transfer
Low hydraulic headless
Operational simplicity
Deep contact basins
Vertical channeling of bubbles
Maintenance of gaskets and piping.
Injector Dissolution
The injector contacting method is commonly used in Europe, Canada, and the United States
(Langlais et al., 1991). Ozone is injected into a water stream under negative pressure, which is
generated in a venturi section, pulling the ozone into the water stream. In many cases, a sidestream of
the total flow is pumped to a higher pressure to increase the available vacuum for ozone injection.
After the ozone is injected into this sidestream, the sidestream containing all the added ozone is
combined with the remainder of the plant flow under high turbulence to enhance dispersion of ozone
into the water. Figure 3-8 illustrates typical in-line and sidestream ozone injection systems.
A. In-line Injector System
Ozone Gas
Influent
Off Gas
Contactor
Injector
B. Sidestream Injector System
Off Gas
Figure 3-8. Sidestream Ozone Injection System
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3. OZONE
The gas to liquid ratio is a key parameter used in the design of injector contacting systems. This ratio
should be less than 0.067 cfm/gpm to optimize ozone transfer efficiency (Langlais et al., 1991).
Meeting this criterion typically requires relatively low ozone dosages and ozone gas concentrations
greater than 6 percent by weight (DeMers and Renner, 1992). High concentration ozone gas can be
generated using a medium-frequency generator and/or liquid oxygen as the feed gas.
To meet the CT disinfection requirements, additional contact time is required after the injector,
typically in a plug flow reactor. The additional contact volume is determined in conjunction with the
applied ozone dosage and estimated residual ozone concentration to satisfy the disinfection CT
requirement.
Table 3-5 summarizes the advantages and disadvantages of injection contacting (Langlais et al.,
1991).
Table 3-5. Injection Contacting Advantages and Disadvantages
Advantages Disadvantages
Injection and static mixing have no moving parts Additional headless (energy usage) due to static mixers
which may require pumping
Very effective ozone transfer Turndown capability limited by injection system
Contactor depth less than bubble diffusion More complex operation and high cost.
Turbine Mixer Contactors
Turbine mixers are used to feed ozone gas into a contactor and mix the ozone with the water in the
contactor. Figure 3-9 illustrates a typical turbine contactor. The illustrated turbine mixer design
shows the motor located outside the basin, allowing for maintenance access. Other designs use a
submerged turbine.
Ozone transfer efficiency for turbine mixers can be in excess of 90 percent. However, the power
required to achieve this efficiency is 2.2 to 2.7 kW-hr of energy per Ib of ozone transferred
(Dimitriou, 1990).
Turbine mixing basins vary in water depth from 6 to 15 ft, and dispersion areas vary from 5 to 15 ft
(Dimitriou, 1990). Again, as with injector contacting, sufficient contact time may not be available in
the turbine basin to meet disinfection CT requirements; consequently additional contact volume may
be required.
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3. OZONE
Contact
Chamber
Off-Gas
Drive Motor
I Ozone-Rich Gas
Ozonated
Water
Figure 3-9. Turbine Mixer Ozone Contactor
Table 3-6 summarizes advantages and disadvantages for the turbine mixer contactor (Langlais et al.,
1991).
Table 3-6. Turbine Mixer Contactor Advantages and Disadvantages
Advantages
Disadvantages
Ozone transfer is enhanced by high turbulence resulting in
small bubble size
Contactor depth less than bubble diffusion
Require energy input
Constant gas flow rate should be maintained, reducing ozone
transfer efficiency
Aspirating turbines can draw off-gas from other chambers for Maintenance requirements for turbine and motor
reuse
Eliminates diffuser clogging concerns
3.2.2.4 Off-gas Destruction Systems
The concentration of ozone in the off-gas from a contactor is usually well above the fatal
concentration. For example, at 90 percent transfer efficiency, a 3 percent ozone feed stream will still
contain 3,000 ppm of ozone in the off-gas. Off-gas is collected and the ozone converted back to
oxygen prior to release to the atmosphere. Ozone is readily destroyed at high temperature (> 350° C
or by a catalyst operating above 100° C) to prevent moisture buildup. The off-gas destruct unit is
designed to reduce the concentration to 0.1 ppm of ozone by volume, the current limit set by OSHA
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3. OZONE
for worker exposure in an eight hour shift. A blower is used on the discharge side of the destruct unit
to pull the air from the contactor, placing the contactor under a slight vacuum to ensure that no ozone
escapes.
3.2.2.5 Instrumentation
Instrumentation should be provided for ozone systems to protect both personnel and the equipment.
Gas phase ozone detectors should be provided in spaces such as generator rooms where ozone gas
may be and personnel are routinely present. An ozone detector is also needed on the outlet from the
off-gas destruct unit to ensure the unit is working properly. These units should be interlocked with
the ozone generator controls to shut down the ozone generation system should excess ozone be
detected. A dew point detector on the feed gas supply just upstream of an ozone generator is
required to protect the generator from moisture in the feed gas (when air is the feed gas). Flow
switches on the cooling water supply are needed to protect the generator from overheating and a
pressure switch to prevent over pressurization.
Other instrumentation can be used to monitor and control the ozone process, although manual control
is adequate for small systems, but most small systems are designed to operate automatically,
particularly in remote areas. Ozone monitors can be used in conjunction with process flow meters to
match ozone dose to process demands and control ozone generation. Sophisticated control schemes
can be implemented to minimize the cost of dosing with ozone and reduce operator attention
requirements. Many systems include residual monitoring at various points in the contactor to
maintain a desired ozone residual and prevent energy-wasting overdosing.
3.2.3 Operation and Maintenance
Even though ozone systems are complex, using highly technical instruments, the process is highly
automated and very reliable, requiring only a modest degree of operator skill and time to operate an
ozone system. Maintenance on generators requires skilled technicians. If trained maintenance staff
are not available at the plant, this work can be done by the equipment manufacturer. Some facilities,
such as the 600 mgd Los Angeles Aqueduct Filtration Plant, use plant mechanics to perform
generator and facilities maintenance. Therefore, backup units are usually installed. Generators
should be checked daily when in operation. After a shutdown, dry air or oxygen should be allowed
to flow through the generator to ensure that any moisture has been purged prior to energizing the
electrodes. At initial start up and after long down times, this process may take up to 12 hours and
usually longer when air is the feed gas. As an alternative, a small flow of dry air can be passed
through the generator continuously when it is in standby mode to maintain the dry condition.
Filters and desiccant in air preparation systems should be changed periodically, with the frequency
depending on the quality of the inlet air and the number of hours in operation. Compressors require
periodic service, depending on the type and operating time. LOX tanks should be periodically
pressure tested. Piping and contact chambers should be inspected periodically to check for leaks and
corrosion.
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Dielectric tubes should be periodically cleaned. This operation should be performed when the
generator efficiency drops 10-15 percent. Cleaning the tubes is a delicate operation as the tubes are
fragile and expensive. Adequate space should be provided for the cleaning operation and for storage
of spare tubes.
3.3 Primary Uses and Points of Application of Ozone
3.3.1 Primary Uses of Ozone
Ozone is used in drinking water treatment for a variety of purposes including:
• Disinfection;
• Inorganic pollutant oxidation, including iron, manganese, and sulfide;
• Organic micropollutant oxidation, including taste and odor compounds, phenolic pollutants, and
some pesticides; and
• Organic macropollutant oxidation, including color removal, increasing the biodegradability of
organic compounds, DBF precursor control, and reduction of chlorine demand.
3.3.1.1 Disinfection
Ozone is a powerful oxidant able to achieve disinfection with less contact time and concentration
than all weaker disinfectants, such as chlorine, chlorine dioxide, and monochloramine (Demers and
Renner, 1992). However, ozone can only be used as a primary disinfectant since it cannot maintain a
residual in the distribution system. Thus, ozone disinfection should be coupled with a secondary
disinfectant, such as chlorine, chloramine, or chlorine dioxide for a complete disinfection system.
3.3.1.2 Iron and Manganese Oxidation
Ozone will oxidize iron and manganese, converting ferrous (2+) iron into the ferric (3+) state and 2+
manganese to the 4+ state. The oxidized forms will precipitate as ferric hydroxide and manganese
hydroxide (AWWA, 1990). The precise chemical composition of the precipitate will depend on the
nature of the water, temperature, and pH. The ozone dose required for oxidation is 0.43 mg/mg iron
and 0.88 mg/mg manganese (Langlais et al., 1991). Iron oxidizes at a pH of 6-9 but manganese is
more effective at a pH of around 8. Also, over-ozonation has no effect on iron, but will resolubilize
manganese, which then should be reduced to manganese dioxide downstream.
3.3.1.3 Oxidation of Taste and Odor Compounds
Ozone is used to oxidize/destroy taste and odor-causing compounds because many of these
compounds are very resistant to oxidation. Suffet et al. (1986) confirmed that ozone is an effective
oxidant for use in taste and odor treatment. They found ozone doses of 2.5 to 2.7 mg/L and 10
minutes of contact time (ozone residual of 0.2 mg/L) significantly reduced taste and odors in the
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specific waters they tested. Most early U.S. water plants (i.e., 1940-1986) installed ozonation
specifically for taste and odor removal.
3.3.1.4 DBP Precursor Control
Early work on oxidation of DBP precursors seemed to indicate that the effects of ozonation, prior to
chlorination, were quite site-specific and unpredictable (Umphries et al., 1979). The key variables
that seem to determine ozone's effect are dose, pH, alkalinity, and, above all, the nature of the ,
organic material. At low pH levels, precursor destruction by ozone is quite effective; however, above
some critical pH, ozone actually is less effective and in fact sometimes increases the amount of
chlorination byproduct precursors. For most humic substances this critical pH is 7.5, which is the
approximate level at which decomposition of ozone to hydroxyl free radicals increases rapidly, thus
increasing organic oxidation rates. Therefore, the implications that at lower pH (approximately 6-7),
at which molecular ozone predominates over the hydroxyl free radical, the initial THM precursor by-
products are different in nature than those formed by the hydroxyl free radicals oxidized at higher pH
levels. This is logical in light of the greater oxidation potential of the hydroxyl free radical over that
of ozone.
As alkalinity increases, it has a beneficial effect on THM formation potential (THMFP) (Langlais et
al., 1991). This is because alkalinity scavenges any hydroxyl free radicals formed during ozonation,
leaving molecular ozone as the sole oxidant, which is only capable of oxidizing organic precursors to
a lower oxidation sequence than does the hydroxyl free radical. Given neutral pH and moderate
levels of bicarbonate alkalinity, THMFP level reductions of 3 to 20 percent have been shown at
ozone doses ranging from 0.2 to 1.6 mg ozone per mg carbon (Singer et al., 1989; Georgeson and
Karimi, 1988).
3.3.1.5 Increase Organic Biodegradation
Ozone can be effective in partially oxidizing organics in the water to biodegradable compounds that
can be removed by biological filtration (Demers and Renner, 1992). This partial oxidation gives rise
to lower molecular weight organics that are more easily biodegradable. This increase in the
biodegradable fraction of organic carbon occurs as a result of moderate to high levels of ozonation.
These ozone levels are typical of the doses commonly applied for disinfection.
3.3.1.6 Coagulation and Filtration Improvement
Ozone has been reported by some to improve coagulation and filtration efficiency (Gurol and
Pidatella, 1993; Farvardin and Collins, 1990; Reckhow et al., 1993; Stolarik and Christie, 1997).
However, others have found no improvement in filter effluent turbidity due to ozonation (Tobiason et
al., 1992; Hiltebrand et al., 1986). Prendiville (1986) collected data from a large water treatment
plant showing that pre-ozonation was more effective than pre-chlorination to reduce filter effluent
turbidities. The cause of the improved coagulation is not clear, but several possibilities have been
offered (Reckhow et al., 1986), including:
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• Oxidation of organic compounds into more polar forms; and
• Oxidation of metals ions to yield insoluble complexes, such as ferric iron complexes.
3.3.2 Points of Application
The typical locations for feeding ozone in a water treatment plant are at the head of the treatment
plant (raw water) pre-ozonation and after sedimentation.
Raw water quality and turbidity and ozone demand (the amount of ozone required for all oxidation
requirements of the water) can be used to assess how to use ozone in the treatment process. Table 3-7
lists the criteria for selecting ozone feed points based on these two parameters. By moving the
ozonation process further downstream after sedimentation, the ozone demand and production of
byproducts are reduced. The advantage of placing ozone ahead of filtration is that biodegradable
organics produced during ozonation can be removed by subsequent biological activity in the filters.
Table 3-7. Criteria for Selecting Ozone Feed Points for Small Systems
Raw Water Quality
Category I
Turbidity < 10 NTU
Ozone Demand < 1mg/L
Category II
Turbidity > 10 NTU
Ozone Demand < 1mg/L
Category III
Turbidity < 10 NTU
Ozone Demand > 1mg/L
Category IV
Turbidity > 10 NTU
Ozone Demand > 1mg/L
Ozone Feed Point(s)
Raw Water or After Sedimentation
After Sedimentation
Raw Water and/or After Sedimentation
After Sedimentation and After First
Stage Filtration, if necessary
Special Considerations
Low ozone demand.
Low disinfection byproducts.
Low biodegradable organics.
Low ozone demand.
High inorganic particulate.
Low biodegradable organics.
High ozone demand
Disinfection byproducts
Biodegradable organics formation
High ozone demand
Disinfection byproducts
Biodegradable organics formation
Source: DsMers and Renner, 1992.
For high quality water with direct filtration, the only practical ozone feed point is the raw water.
Category II (Table 3-7) water is characterized by low ozone demand and high turbidity. This water
quality indicates the presence of inorganic material, such as clay or silt particles. For ozone to be
most effective for Category II water disinfection, it should be added after either pre-sedimentation or
conventional sedimentation.
Raw water with low turbidity and high ozone demand (Category III, Table 3-7) contains dissolved
constituents, not suspended, that contribute to a high ozone demand. An example of this type of
water is a ground water containing bromide ion, iron, manganese, color, or organics. For this water
quality, ozone can be added to either the raw water or after sedimentation. If the water contains
organic constituents that become more biodegradable by ozonation, a biological treatment step (see
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3. OZONE
Section 3.3.4) may be required. The presence of oxidizable organic constituents or bromide ion will
generate disinfection byproducts upon ozonation.
Category IV (Table 3-7) water would be considered the most difficult water to treat with ozone due
to its high turbidity and high ozone demand. An example of this water quality is surface water
containing high concentrations of organic material and inorganic particles. The most effective use of
ozone for this water quality is after sedimentation and possibly after filtration. If the water has an
extremely high ozone demand, dual ozone feed points may be required to achieve disinfection goals,
because the presence of large amounts of organic material may require a biological treatment step
and may generate disinfection byproducts.
3.3.3 Impact on Other Treatment Processes
Ozonation does have an impact oil other processes at the water treatment facility. The impacts of
ozone addition include the following:
• The use of ozone generates biodegradable organic matter (BOM) that can result in biological
growth which may also increase corrosion rates in distribution systems if not removed by
biologically active filtration. When ozonation is placed before biological filters, it can impact the
filters by increasing biological growth and increasing backwash frequency.
• Ozone is a strong oxidant that reacts with other oxidants, such as chlorine, chlorine dioxide, and
monochloramine.
• Ozone oxidation of iron and manganese generates insoluble oxides that should be removed by
sedimentation or filtration. These insoluble oxides also impact the filters by increasing load on
the filters and increasing backwash frequency.
• Using pre- and/or internal ozone on most raw waters reduces the subsequent chlorine, chlorine
dioxide, or monochloramine demand of the finished water so as to allow a stable chlorine-
compound residual to be maintained at a much lower level.
The reader is referred to EPA's Simultaneous Compliance Guidance Document (expected to be
available in 1999) for additional information regarding the interaction between oxidants and other
treatment processes.
3.3.4 Biologically Active Filtration
Ozonation typically increases the biodegradability of NOM in water because many large organic
molecules are converted into smaller organic molecules that are readily biodegradable. This increase
in biodegradable dissolved organic carbon (BDOC) can lead to accelerated bacterial growth and
regrowth in the distribution system if not removed in the treatment plant. LeChevallier et al. (1992)
found that AOC levels less than 100 ppb may be necessary to control excessive bacterial regrowth in
the distribution system if not removed in the treatment plant.
When ozonation is placed upstream of filtration, and environmental conditions such as dissolved
oxygen, pH, and temperature are favorable, microbiological activity is increased in the filter and
BDOC/AOC removal is enhanced. Ozone addition not only increases the biodegradability of the
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dissolved organics, but also introduces large amounts of oxygen to the water, thus, creating an
excellent environment for biological growth on the filter media. The advantages of biologically
active filtration (Price, 1994) include the following which are all being met in most U.S. plants using
ozone:
• Production of a biologically stable water that does not promote excessive bacterial growth and
regrowth in the distribution system.
• Removal of NOM that can serve as precursors to byproduct formation as a result of residual
disinfection with free or combined chlorine.
• Ozone oxidation as a primary disinfectant prior to biologically active filtration reduces the
BDOC concentration in finished water, thus reducing chances of regrowth.
• Reduction of the residual disinfectant demand of the product water so that proposed regulations
limiting the maximum disinfectant residual can be met.
• Removal or control of ozonation byproducts.
Biological activity can be supported on slow sand, rapid rate, and GAC media because these media
provide a surface for bacteria to attach. Factors such as available surface area, hydraulic loading rate,
contact time, availability of nutrients, temperature, and others will determine the performance and
BDOC removal efficiency. Biomass develops to higher levels on GAC because of the rougher
surface characteristics than on anthracite and sand.
3.3.4.1 Slow Sand Filters
Ozone addition prior to slow sand filtration can increase the efficiency of TOC removal by about 35
percent (Rachwal et al., 1988; Zabel, 1985). Ozone addition can also increase the efficiency of
BDOC removal with slow sand filters (Eighmy et al., 1991; Malley et al., 1993).
3.3.4.2 Rapid Rate Filters
Research in the area of biologically active rapid rate filters has focused on the reduction of
assimilable organic carbon (AOC) instead of BDOC. While studies have shown rapid rate filtration,
employing either sand or dual media lowers AOC levels following ozonation, AOC does not measure
all the BDOC. AOC measures only that portion of the BDOC that is more easily assimilable or more
easily biodegradable under specific laboratory conditions by two specific microorganisms. Research
data shows that biodegradation of AOC can occur in rapid rate filters. The data should be viewed
with caution, since the more slowly biodegradable DOC, not measured by AOC, may be passing
through rapid rate filters.
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3.3.4.3 Granular Activated Carbon
GAC is made biologically active by the deliberate introduction of sufficient dissolved oxygen to
water just before passing through GAC columns (Katz, 1980). The high surface area and long
retention time in GAC provide an ideal environment to enhance biological growth.
Although ozone actually increases the amount of BDOC, the efficiency of subsequent biodegradation
on GAC can be such that the BDOC in GAC effluent is lower than the BDOC in the ozone influent
(Langlais et al., 1991). The degree to which biodegradable DOC is removed by ozone/GAC depends
upon the process conditions of temperature, amount of BDOC, and the GAC column loading rate,
measured by empty bed contact time (EBCT). For example, with an influent BDOC of 0.65 mg C/L
and a 10 minute EBCT, one would expect an effluent BDOC of 0.25 mg C/L. The effluent BDOC
then could be lowered by either:
• Adding ozone, which would increase the GAC influent BDOC and, therefore, lower the
effluent BDOC; or
• Adding more GAC or decreasing the loading rate, which would extend the EBCT and lower
the effluent BDOC (Billen et al., 1985, as cited in Langlais et al., 1991).
Huck et al. (1991) reported results from AOC profiles measured in a pilot treatment plant. The plant
treated Saskatchewan River water and included coagulation, flocculation, and sedimentation prior to
ozonation. Following ozonation, the water was filtered through a dual media (anthracite-sand) filter
followed by GAC adsorption. The results demonstrate:
• Variable AOC removal through coagulation, flocculation, and sedimentation (80 percent to
zero);
• Increased AOC after ozonation;
• AOC removal through dual media filtration improving at lower hydraulic loading rates and
filtered effluent AOC often less than raw water AOC, but highly variable; and
• AOC levels after GAC were low, almost always below raw water AOC concentrations and
adsorption appears to contribute to some immediate AOC removal.
3.3.5 Pathogen Inactivation and Disinfection Efficacy
Ozone has a high germicidal effectiveness against a wide range of pathogenic organisms including
bacteria, protozoa, and viruses. Because of its high germicidal efficiency, ozone can be used to meet
high inactivation required by water treatment systems with or without filters. However, ozone
cannot be used as a secondary disinfectant because the ozone residual decays too rapidly. The ozone
disinfection efficiency is not affected by pH (Morris, 1975), although because of hydroxyl free
radicals and rapid decay, efficiency is the same but more ozone should be applied at high pH to
maintain "C".
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3.3.6 Inactivation Mechanisms
Inactivation of bacteria by ozone is attributed to an oxidation reaction (Bringmann, 1954; Chang,
1971). The first site to be attacked appears to be the bacterial membrane (Giese and Christensen,
1954) either through the glycoproteins or glycolipids (Scott and Lesher, 1963) or through certain
amino acids such as typtophan (Goldstein and McDonagh, 1975). In addition, ozone disrupts
enzymatic activity of bacteria by acting on the sulfhydryl groups of certain enzymes. Beyond the cell
membrane and cell wall, ozone may act on the nuclear material within the cell. Ozone has been
found to affect both purines and pyrimidines in nucleic acids (Giese and Christensen, 1954; Scott and
Lesher, 1963).
The first site of action for virus inactivation is the virion capsid, particularly its proteins (Cronholm et
al., 1976 and Riesser et al., 1976). Ozone appears to modify the viral capsid sites that the virion uses
to fix on the cell surfaces. High concentrations of ozone dissociate the capsid completely. One
researcher found that the mechanism of ozone inactivation of bacteriophage f2 ribonucleic acid
(RNA) included releasing RNA from the phage particles after the phage coat was broken into many
pieces (Kim et al., 1980). This finding suggests that ozone breaks the protein capsid, thereby
liberating RNA and disrupting adsorption to the host pili. Further, the naked RNA may be
secondarily inactivated by ozone at a rate less than that for RNA within the intact phage. The
mechanism for inactivation of deoxyribonucleic acid (DNA) bacteriophage T4 has been found to be
quite similar to RNA inactivation: ozone attacks the protein capsid, liberates the nucleic acid, and
inactivates the DNA (Sproul et al., 1982). In contrast, more recent work on the tobacco mosaic virus
(TMV) shows that ozone has a specific effect on RNA. Ozone was found to attack both the protein
coat and RNA. The damaged RNA cross-links with amino acids of the coat protein subunits. The
authors concluded that TMV loses its infectivity because of its loss of protein coating.
Microscopic observation of inactivation of trophozoites ofNaegleria and Acanthamoeba showed that
they were rapidly destroyed and the cell membrane was ruptured (Perrine et al., 1984). Perrine and
Langlais showed that ozone affect the plugs in Naegleria gruberi cysts (Langlais and Perrine, 1984).
Depending on the ozonation conditions, these plugs were completely removed or were partially
destroyed. It has been speculated that ozone initially affects the Giardia muris cysts wall and makes
it more permeable (Wickramanayake, 1984c). Subsequently, aqueous ozone penetrates into the cyst
and damages the plasma membranes, additional penetration of ozone eventually affects the nucleus,
ribosomes, and other ultrastructural components.
3.3.7 Disinfection Parameters
Hoigne" and Bader demonstrated that the rate of decomposition of ozone is a complex function of
temperature, pH, and concentration of organic solutes and inorganic constituents (Hoigne and Bader,
1975, and 1976). The following sections describe the effects that pH, temperature, and suspended
matter have on the reaction rate of ozone and pathogen inactivation.
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The ability to maintain a high aqueous ozone concentration is critical from a regulatory disinfection
compliance standpoint. This means that factors that accelerate ozone decomposition are undesirable
for inactivation because the ozone residual dissipates faster and therefore reduces the CT credit,
requiring a corresponding increase in the ozone applied, thus increasing cost.
3.3.7.7 pH
Studies have indicated that pH has little effect on the ability of dissolved ozone residuals to inactivate
acid-fast bacteria, such as Mycobacteria and Actinomycetes (Farooq, 1976). A slight decrease has
been found in the virucidal efficacy of ozone residuals as pH decreased (Roy, 1979). However, the
opposite effect was observed by Vaughn et al. (1987) (cited in Hoff, 1986). Changes in disinfection
efficacy with variations in pH appear to be caused by the ozone decomposition rate. Ozone
decomposition occurs faster in higher-pH aqueous solutions and forms various types of oxidants with
differing reactivities (Langlais et al., 1991). Tests carried out at constant ozone residual
concentration and different pH values showed that the degree of microorganism inactivation
remained virtually unchanged (Farooq et al., 1977). More recent studies have indicated decreased
virus inactivation by ozone at alkaline pH (pH 8 to 9) for poliovirus 1 (Harakeh and Butler, 1984)
and rotaviruses SA-11 and Wa (Vaughn et al., 1987).
Inactivation of Giardia murls cysts was found to improve when the pH increased from 7 to 9
(Wickramanayake, 1984a). This phenomenon was attributed to the possible changes in cyst
chemistry making it easier for ozone to react with the cyst constituents at the higher pH levels.
However, the same study found that inactivation ofNaegleria gruberi cysts was slower at a pH 9
than at lower pH levels, thereby indicating that pH effects are organism-specific.
3.3.7.2 Temperature
As temperature increases, ozone becomes less soluble and less stable in water (Katzenelson et al.,
1974); however, the disinfection and chemical oxidation rates remain relatively stable. Studies have
shown that although increasing the temperature from 0 to 30°C can significantly reduce the solubility
of ozone and increases its decomposition rate, temperature has virtually no effect on the disinfection
rate of bacteria (Kinman, 1975). In other words, the disinfection rate was found to be relatively
independent of temperature at typical water treatment plant operating temperatures despite the
reduction in solubility and stability at higher temperatures.
3.3.7.3 Suspended Matter
Ozone inactivation of viruses and bacteria contained in aluminum floe (in the size range comparable
to those that could typically escape filtration) was not reduced at floe turbidity levels of 1 and 5 NTU
(Walsh et al., 1980). This study demonstrated that the microorganisms received no protection from
the aluminum floe. Similar results have been obtained for poliovirus 1, coxsackie virus A9, and E.
coli associated with bentonite clay (Boyce et al., 1981). However, adsorption of the f2 bacteriophage
at 1 and 5 NTU of bentonite clay was found to retard the rate of inactivation of ozone (Boyce et al.,
1981).
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In some instances, river waters heavily polluted with organic matter were ozonated, and the results
indicated a degradation of large organic molecules into fragments more easily metabolized by
microorganisms. This fragmentation coupled with the inability of ozone to maintain an active
concentration in the distribution system, has led to increased slime growth and, consequently, water
quality deterioration during distribution (Troyan and Hanson, 1989).
3.3.8 inactivation of Microorganisms
The following sections contain a description of the disinfection efficiency of ozone in terms of
bacteria, virus, and protozoa inactivation.
3.3.8.1 Bacterial Inactivation
Ozone is very effective against bacteria. Studies have shown the effect of small concentrations of
dissolved ozone (i.e., 0.6 jag/L) on E. coli. (Wuhrmann and Meyrath, 1955) and Legionella
pneiunophila (Domingue, et al., 1988). E. coli. levels were reduced by 4 logs (99.99 percent
removal) in less than 1 minute with a ozone residual of 9 jag/L at a temperature of 12°C. Legionella
pnewnophila levels were reduced by greater than 2 logs (99 percent removal) within a minimum
contact time of 5 minutes at a ozone concentration of 0.21 mg/L. Results similar to those obtained
for E. coli. have been found for Staphyloccus sp. and Pseudomonas fluorescens inactivation.
Streptococcus faecalis required a contact time twice as long with the same dissolved ozone
concentration, and Mycobacterium tuberculosis required a contact time six times as long for the same
reduction level as E. coli.
In regard to vegetative bacteria, E. coli is one of the most sensitive types of bacteria. Furthermore,
Significant difference has been found among all the Gram-negative bacillae, including E. coli and
other pathogens such as Salmonella, which are all sensitive to ozone inactivation. whereas the Gram-
positive cocci (Staphyloccus and Streptococcus), the Gram-positive bacillae (Bacillus), and the
Mycobacteria are the most resistant forms of bacteria. Sporular bacteria forms are always far more
resistant to ozone disinfection than vegetative forms (Bablon, et al., 1991), but all are easily
destroyed by relatively low levels of ozone.
3.3.8.2 Protozoa Inactivation
Protozoan cysts are much more resistant to ozone and other disinfectants than vegetative forms of
bacteria and viruses. Giardia lamblia has a sensitivity to ozone that is similar to the sporular forms
of Mycobacteria. Both Naegleria and Acanthamoeba cysts are much more resistant to ozone (and all
other disinfectants) than Giardia cysts. (Bablon et al., 1991). CT products for 99 percent inactivation
of Giardia lamblia and N. gruberi at 5°C were 0.53 and 4.23 mg • min/L, respectively
(Wickramanayake et al., 1984a and 1984b). Data available for inactivation of Cryptosporidium
oocysts, suggest that among protozoans, this microorganism is more resistant to ozone (Peeters et al.,
1989; Langlais et al., 1990). One study found that Cryptosporidium oocysts are approximately 10
times more resistant to ozone than Giardia (Owens et al., 1994).
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3.3.8.3 Virus Inactivation
Typically, viruses are more resistant to ozone than vegetative bacteria but less resistant than sporular
forms of Mycobacteria (Bablon, et al., 1991). The most sensitive forms of viruses are phages, and
there seems to be little difference between the polio- and coxsackie viruses. The sensitivity of human
rotavirus to ozone was determined to be comparable to that of Mycobacteria and polio- and
coxsackie viruses (Vaughn et al., 1987).
Keller et al. (1974) studied ozone inactivation of viruses by using both batch tests and pilot plant
data. Inactivation of poliovirus 2 and coxsackie virus B3 was more than 3 logs (99.9 percent) in the
batch tests with an ozone residual of 0.8 mg/L and 1.7 mg/L and a contact time of 5 minutes. Greater
than 5 log (99.999 percent) removal of coxsackie virus was achieved in the pilot plant with an ozone
dosage of 1.45 mg/L, which provided an ozone residual of 0.28 mg/L in lake water.
3.3.8.4 CT Curves for Giardia Lamblia
CT values shown in Figure 3-9 are based on disinfection studies using in vitro excystation of Giardia
lamblia. CT values obtained at 5°C and pH 7 were used as the basis for deriving the CT values at
other temperatures. A safety factor of 2 has been applied to the values shown in Figure 3-9.
2.0
1.8
-*—0.5-log Inactivation
-*— 1-log Inactivation
-*— 1.5-log Inactivation
••••&•••• 2-log Inactivation
—•— 2.5-log Inactivation
-8— 3-log Inactivation
15
Temperature (°C)
Figure 3-9. CT Values for Inactivation of Giardia Cysts by Ozone (pH 6 to 9)
April 1999
3-25
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CT values shown in Figure 3-10 for achieving 2-log inactivation of viruses were determined by
applying a safety factor of 3 to data obtained from a previous study on poliovirus 1 (Roy et al.,
1982). CT values for 3 and 4-log removal were derived by applying first order kinetics and assuming
the same safety factor of 3. Data obtained at a pH of 7.2 was assumed to apply for the pH range of 6
to 9.
Several research groups have investigated the efficiency of ozone for Cryptosporidium oocyst
inactivation. Table 3-8 summarizes CT values obtained for 99 percent inactivation of
Cryptosporidium oocysts. Results indicate that ozone is one of the most effective disinfectants for
controlling Cryptosporidium (Finch, et al., 1994) and that Cryptosporidium muris may be slightly
more resistant to ozonation than Cryptosporidium parvum (Owens et al., 1994). A wide range of CT
values has been reported for the same inactivation level, primarily because of the different methods
of Cryptosporidium measurements employed and pH, temperature, and above all, ozonation
conditions.
1.2
—*-2-log Inactivation
-*— 3-log Inactiviation
~*~ 4-log Inactivation
Temperature (°C)
Figure 3-10. CT Values for Inactivation of Viruses by Ozone (pH 6 to B)
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Table 3-8. Summary of Reported Ozonation Requirements for 99 Percent
Inactivation of Cryptosporidium Oocycts
Species
C. baileyi
C. muris
C. pan/urn
C. parvum
C. parvum
C. parvum
Ozone
protocol
Batch liquid,
modified batch
ozone
Flow through
contactor,
continuous gas
Batch liquid,
batch ozone
Batch liquid,
batch ozone
Batch liquid,
continuous gas
Flow through
contactor,
continuous gas
Ozone
residual
(mg/L)
0.6 & 0.8
0.50
0.50
0.77
0.51
1.0
Contact time
(min)
4
18
7.8
6
8
5&10
Temperature
(•C)
• 25
22-25
7
22
Room
25
22-25
CT
(mg.min/L)
2.4-3.2
7.8
9.0
3.9
4.6
4
5-10
5.5
Reference
Langlais et al.,
1990
Owens et al.,
1994
Finch et al.,
1993
Peelers et al.,
1989
Korich et al.,
1990
Owens et al.,
1994
Ozone dose and contact time (CT) requirements for the inactivation of Cryptosporidium oocysts in
drinking water when using ozone has not been established similar to the CT values for viruses and
Giardia cyst inactivation. Inactivation requirements (log removals) for Cryptosporidium oocysts
have not been established. In addition, as shown in Table 3-8, the CT requirements reported in the
literature vary from study to study which adds uncertainty to design CT requirements for specific
applications or regulatory needs.
3.4 Ozonation Disinfection Byproducts
Ozone does not form halogenated DBFs (TTHMs and HAASs) when participating in
oxidation/reduction reactions with NOM but it does form a variety of organic and inorganic
byproducts. Table 3-9 and Figure 3-11 show the principal known byproducts associated with
ozonation. However, if bromide ion is present in the raw water halogenated DBFs may be formed.
These brominated DBFs appear to pose a greater health risk than non-brominated DBFs.
April 1999
3-27
EPA Guidance Manual
Alternative Disinfectants and Oxidants
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a OZONE
Table 3-9. Principal Known Byproducts of Ozonation
Disinfectant Byproducts
Aldehydes
Formaldehyde
Acetaldehyde
Glyoxal
Methyl Glyoxal
Acids
Aldo- and Ketoacids
Pyruvic acid
Brominated Byproducts*
Bromate ion
Bromoform
Brominated acetic acids
Oxalic acid
Succinic acid
Formic acid
Acetic acid
Bromopicrin
Brominated acetonitriles
Others
Hydrogen peroxide
•Bromlnated byproducts are produced only in waters containing bromide ion
Source: Singer, 1992.
Although ozone is an effective oxidant and disinfectant, it should not be relied upon as a secondary
disinfectant to maintain a residual in the distribution system. Monochloramine is attractive for this
purpose because it produces little to no halogenated DBFs. Chlorine is a candidate for secondary
disinfectant but the ozonated water may actually produce either more or less DBFs following the
addition of free chlorine depending on the nature of the organic material following ozonation unless
biologically active filtration precedes the addition of chlorine. The principal benefit of using ozone
for controlling DBF formation is that it allows free chlorine to be applied later in the treatment
process after precursors have been removed and at lower doses, thereby reducing DBPFP.
HOBr •<—>- H+ + OBr
pKa = 8.7 @ 25°C
Organic Precursors Aldehydes and
Oxidized Organics
NH2Br, NHBr2
Brominated
Organic DBPs
Organic
Precursors
+ OBr"
Figure 3-11. Principal Reactions Producing Ozone Byproducts
Application of a secondary disinfectant following ozonation requires special consideration for
potential interaction between disinfectants. For example, chloral hydrate formation has been
EPA Guidance Manual
Alternative Disinfectants and Oxidants
3-28
April 1999
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, ^ 3. OZONE
observed when using chlorine as a secondary disinfectant with ozone (McKnight and Reckhow,
1992; Logsdon et al., 1992). The ozonation byproduct of acetaldehyde is a known precursor for
chloral hydrate, a byproduct of chlorination. Enhancement of chloral hydrate has not been observed
when monochloramine is applied as the secondary disinfectant, or if biologically active filtration is
used following ozonation and prior to chlorination (Singer, 1992). Chloropicrin formation from free
chlorine appears to be enhanced by pre-ozonation (Hoigne and Bader, 1988) in the absence of
biologically active filtration prior to addition of chlorine.
Byproducts such as aldehydes, ketones, acids, and others will be formed upon ozonation of water.
The primary aldehydes that have been measured are: formaldehyde, acetaldehyde, glyoxal, and
methyl glyoxal (Glaze et al., 1991). Total aldehyde concentration in drinking water disinfected with
ozone range from less than 5 p.g/L to 300 jag/L, depending on the TOC concentration and the applied
ozone to organic carbon ratio (Van Hoof et al., 1985; Yamada and Somiya, 1989; Glaze et al., 1989a;
Krasner et al., 1989; Glaze et al., 1991; LeLacheur et al., 1991). Aldehydes with higher molecular
weights have also been reported (Glaze et al., 1989b). Other organic byproducts of ozonation are
indicated in Table 3-9
Ozonation of a source water containing bromide ion can produce brominated byproducts, the
brominated analogues of the chlorinated DBFs. Song et al. (1997) found that bromate ion formation
is an important consideration for waters containing more than 0.10 mg/L bromide ion. These
brominated byproducts include bromate ion, bromoform, the brominated acetic acids and
acetonitriles, bromopicrin, and cyanogen bromide (if ammonia is present). An ozone dose of 2 mg/L
produced 53 f^g/L of bromoform and 17 ng/L of dibromoacetic acid in a water containing 2 mg/L of
bromide ion (McGuire et al., 1990). Ozonation of the same water spiked with 2 mg/L bromide ion
showed cyanogen bromide formation of 10 |ig/L (McGuire et al., 1990). Furthermore, ozone may
react with the hypobromite ion to form bromate ion (Amy and Siddiqui, 1991; Krasner et al., 1993), a
probable human carcinogen (Regli et al., 1992). Bromate ion concentrations in ozonated water up to
60 jig/L have been reported (Amy and Siddiqui, 1991; Krasner et al., 1993). Note that the amount of
bromide ion incorporated into the measured DBFs accounts for only one-third of the total raw water
bromide ion concentration. This indicates that other brominated DBFs exist that are not yet
identified (Krasner et al., 1989; MWDSC and JMM, 1992). Figure 3-12 shows the major pathways
for bromate ion formation.
April 1999 3-29 " EPA Guidance Manual
Alternative Disinfectants and Oxidants
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3. OZONE
HOBr
>
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o;
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Source: Song el al., 1997.
Figure 3-12. Main Pathways of Bromate Ion Formation when Ozone Reacts
with Bromide Ion
3.4.1 Ozone Byproduct Control
The primary factors affecting the speciation and concentrations of brominated byproducts are pH and
the ozone-to-bromide ion and TOC-to-bromide ion ratios (Singer, 1992). Bromate ion formation can
be controlled by ozonation at acidic pH values of which hypobromous acid dominates over the now
absent hypobromite ion (Haag and Hoign6, 1984; Amy and Siddiqui, 1991; Krasner et al., 1993).
Conversely, under alkaline pH conditions, ozone can oxidize the hypobromous acid further to
produce bromate ion. At low pH values, the brominated organic byproducts are favored, while at
greater pH values, bromate ion formation is favored. Therefore, the application of ozone may be
limited for source waters containing bromide ion. Bromate ion formation can be controlled by
lowering the ambient bromide ion concentration, lowering the ozone residual, and lowering the
ozonation pH. The addition of ammonia with ozonation to form bromamines reduces the formation
of both bromate ion and organic byproducts (Amy and Siddiqui, 1991; MWDSC and JMM, 1992).
However, ammonia can act as a nutrient for nitrifying bacteria.
EPA Guidance Manual
Alternative Disinfectants and Oxidants
3-30
April 1999
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3. OZONE
The organic acid and aldehyde byproducts of ozonation discussed above appear to be readily
biodegradable and are a component of the assimilable organic carbon (AOC) or biodegradable
organic carbon (BDOC). Ozonation increases the BOM by oxidation. Therefore, if water disinfected
with ozone is coupled with a biologically active process (i.e., biological active carbon), removal of
these biodegradable byproducts can be reduced. The use of biologically active filters, maintained by
discontinuing the application of a disinfectant to the filters, has been shown to successfully remove
aldehydes and other compounds representing a portion of the BDOC in a water (Bablon et al., 1988;
Rittman, 1990; Reckhow et al., 1992). See Section 3.3.4 for a detailed discussion on biological
active filters.
A recent study has shown that bromate ion and brominated organics can be controlled during
ozonation by the following techniques (Song et al., 1997):
• Low pH decreases bromate ion formation while increasing brominated organic formation;
• Ammonia addition with short ozone contact time decreases both bromate ion and brominated
organic formation;
• Hydrogen peroxide decreases brominated organic formation and may increase or decrease
bromate ion formation, depending on other water quality parameters; and
• Low ozone DOC ratio leads to low bromate ion and brominated organic formation.
3.5 Status of Analytical Methods
During operation of an ozonation system it is necessary to analyze for ozone in both the liquid and
gas phase to determine the applied ozone dose, ozone transfer efficiency and (for primary
disinfection) residual ozone level. The gas stream exiting the ozone generator is monitored for ozone
content to determine the applied ozone dose. The off-gas exiting the ozone contactor is monitored to
determine the amount of ozone transferred to the liquid phase in the contactor and to calculate the
ozone transfer efficiency. The disinfected water exiting the ozone contactor is monitored for residual
ozone to ensure that CT values are met.
In addition, the ambient air in any ozone generating or handling room and ozone destruct off-gas are
monitored for ozone concentration to protect workers in the event of leakages or destruct system
failure.
3.5.1 Monitoring of Gas Phase Ozone
Points in an ozone treatment system where gas-phase ozone is monitored include:
• Ozone generator output;
• Contactor off-gas;
• Ozone destruct off-gas; and
April 1999 3-31 EPA Guidance Manual
Alternative Disinfectants and Oxidants
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3. OZONE
• Ambient air in ozone process areas.
The range of ozone concentrations to be measured in the gas phase varies from less than 0.1 ppm by
volume (0.2 mg/m3 NTP) in ambient air and ozone destruct off-gas, to 1 to 2 percent (10 g/m3 NTP)
in contactor off-gas, to as high as 15 percent by weight (200 g/m3 NTP) in the ozone generator
output.
Analytical methods for monitoring gas-phase ozone include:
• UV absorption;
• lodometric methods;
• Chemiluminescence; and
• Gas-Phase titration.
Table 3-10 presents the working range, expected accuracy and precision, operator skill level
required, interferences, and current status for gas phase ozone analysis.
3.5.1.1 UV Absorption
Gaseous ozone absorbs light in the short UV wavelength region with a maximum absorbance at
253.7 nm (Gordon et al., 1992). Instruments for measuring ozone by the absorption of UV radiation
are supplied by several manufacturers for gas concentrations below 0.5 ppm by volume (1 g/m3
NTP). In general, these instruments measure the amount of light absorptions when no ozone is
present and the amount of light absorptions when ozone is present. The meter output is the difference
of the two readings, which is directly related to the actual amount of ozone present. The International
Ozone Association (IOA) has accepted this procedure (IOA, 1989).
EPA Guidance Manual 3-32 April 1999
Alternative Disinfectants and Oxidants
-------
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3.5.1.2 lodometric Methods
lodometric procedures have been used for all of the ozone concentration ranges encountered in water
treatment plants (Gordon et al., 1992). This includes measurement of ozone directly from the
generator and of ozone as stripped from aqueous solution. For the iodometric method, the ozone-
containing gas is passed into an aqueous solution containing excess potassium iodide.
All other oxidizing materials act as interferences with iodometry (Gordon et al., 1992). Nitrogen
oxides (that may be present when ozone is generated in air) also act as interferences to iodometric
methods. The effects of nitrogen oxides may be eliminated by passing the ozone-containing gas
through absorbents such as potassium permanganate that are specific for nitrogen oxide gases.
However, no iodometric method is recommended for the determination of ozone in solution because
of the unreliability of the method (Gordon et al., 1989).
3.5.1.3 Chemiluminescence
Chemiluminescence methods can be used for the determination of low concentrations of ozone in the
gas phase (Gordon et al., 1992). One of the most commonly used methods is ethylene
Chemiluminescence. Gas-phase ozone can be measured using the chemiluminescent reaction between
ethylene and ozone. This method is specific to ozone and is suitable for measurement of ozone in the
ambient air. The ethylene Chemiluminescence procedure was adopted in 1985 by the EPA as its
reference method for determining ozone in the ambient atmosphere (McKee et al., 1975).
Chemiluminescent instruments are approved by the EPA for monitoring ambient ozone
concentrations of 0 to 0.5 or 0 to 1.0 ppm by volume. With regular calibration, this type of
instrument is capable of providing reliable analysis of any ozone in the ambient air from an
ozonation plant.
An alternative to ethylene Chemiluminescence is rhodamine B/gallic acid Chemiluminescence, which
avoids the handling of ethylene (Gordon et al., 1992). This alternative method is considerably more
complex than the more common ethylene Chemiluminescence instruments. The sensitivity of this
method tends to drift and a procedure has been developed by which corrections are made for the
sensitivity on a frequent basis (Van Dijk and Falkenberg, 1977). Given the wide availability of
ethylene Chemiluminescence monitors and their approval by EPA, ethylene monitors should be
considered before rhodamine B/gallic acid monitors.
3.5.1.4 Gas-Phase Titration
Two gas-phase titration methods have been studied as possible calibration methods for ambient
ozone analyzers and monitors (Gordon et al., 1992). These procedures are based on titration with
nitric oxide and back titration of excess nitric oxide (Rehme et al., 1980). These gas phase titration
procedures, evaluated by EPA, were compared with UV absorption and iodometry as calibration
methods for ethylene chemiluminescent ambient air ozone analyzers. As a result of these
comparisons, UV absorption has been specified as the method of calibration for ambient ozone
EPA Guidance Manual 3-34 April 1999
AttemaUve Disinfectants and Oxidants
-------
3. Ozcwe
analyzers. Therefore, gas-phase titration methods are not recommended for use at ozonation facilities
(Gordon et al., 1992).
3.5.2 Monitoring of Liquid Phase Residual Ozone
There are numerous methods for monitoring ozone in aqueous solutions. Gordon et al. (1992)
recommended the following methods for analyzing residual ozone:
• Indigo colorimetric method;
• Acid chrome violet K (ACVK) method;
• Bis-(Terpyridine) iron(II) method; and
• Stripping into the gas-phase.
Table 3-11 shows the working range, expected accuracy and precision, operator skill level required,
interferences and current status for liquid phase residual ozone analytical methods.
3.5.2.1 Indigo Colorimetric Method
The indigo colorimetric method is the only method for monitoring residual ozone in Standard
Methods, 1995. The indigo colorimetric method is sensitive, precise, fast, and more selective for
ozone than other methods. There are two indigo colorimetric methods: spectrophotometric and
visual. For the spectrophotometric procedure the lower limit of detection is 2 }J.g/L, while for the
visual procedure the detection limit is 10 fig/L.
Hydrogen peroxide, chlorine, manganese ions, ozone decomposition products, and the products of
organic ozonation exhibit less interference with the indigo colorimetric method than any of the other
methods (Langlais et al., 1991). However, the masking of chlorine in the presence of ozone can make
the indigo method problematic. In the presence of hypobromous acid, which forms during ozonation
of bromide-ion containing methods, an accurate measurement cannot be made with this method
(Standard Methods, 1995).
April 1999 3-35 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
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3.5.2.2 ACVK Method
ACVK is readily bleached by ozone, which serves as the basis for this spectrophotometric procedure
first reported by Masschelein (1977). The procedure has been developed further for the measurement
of ozone in the presence of chlorine (Ward and Larder, 1973).
The advantages of this method (Langlais et al., 1991) are:
• Its ease of execution and stability of the dye reagent solution;
• A linear calibration curve over the range of 0.05 to 1.0 mg/L of ozone; and
• The apparent lack of interferences from low levels of manganese, chlorine or combined
chlorine (up to 10 mg/L), organic peroxides, and other organic oxidation products.
3.5.2.3 Bis-(Terpyridine)lron(ll) Method
Bis-(Terpyridine)Iron(II) in dilute hydrochloric acid solution reacts with ozone to change the
absorbance spectra measured at 552 nm (Tomiyasu and Gordon, 1984). The only known interferent
is chlorine. However, chlorine interferences can be masked by the addition of malonic acid.
Alternatively, since the reaction between chlorine and the reagent is slower than the reaction with
ozone, the spectrophotometric measurements can be carried out immediately after reagent mixing to
reduce the chlorine interference. Chlorine dioxide does not interfere (Gordon et al, 1992).
The primary advantages of the Bis-(Terpyridine)Iron(II) Method are its lack of interferences, low
limits of detection (4 u.g/L), broad working range (up to 20 mg/L), excellent reproducibility, and
agreement with the indigo colorimetric method.
3.5.2.4 Stripping into the Gas-Phase
In this indirect procedure, residual ozone is stripped from the solution using an inert gas. The amount
of ozone present in the gas phase is then analyzed by gas-phase analytical methods such as UV
absorption or chemiluminescence described previously. This stripping technique was developed to
minimize the complications caused by the presence of other oxidants in solution.
The success of this procedure initially depends upon the ability to strip ozone from the treated water
without any decomposition. Since stripping conditions such as temperature, pH, and salinity can
vary, the reliability of this method is suspect (Langlais et al., 1991).
3.5.3 Bromate Monitoring for Systems Using Ozone
The DBPR requires that community water systems and non-transient non-community water systems that
use ozone for disinfection or oxidation must monitor their system for bromate. These systems are
required to take one sample per month for each treatment plant with samples collected at the entrance to
the distribution system while the ozonation system is operating under normal conditions.
April 1999 3-37 . EPA Guidance Manual
Alternative Disinfectants and Oxidants
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3. OZONE
The DBPR provides reduced monitoring opportunities (i.e., quarterly rather than monthly samples) if the
system demonstrates that the average source water bromide concentration is less than 0.05 mg/L based
upon representative monthly bromide measurements for one year. Systems can remain on the reduced
monitoring schedule until the running annual average source water bromide concentration, computed
quarterly, is equal to or greater than 0.05 mg/L based upon representative monthly measurements.
For compliance monitoring for bromate, systems must use the ion chromatography analytical method as
specified in USEPA Method 300.1, Determination of Inorganic Anions in Drinking Water by Ion
Chromatography, Revision 1.0 (USEPA, 1997).
If the average of samples covering any consecutive four-quarter period exceeds the MCL, the system is in
violation of the MCL and must notify the public pursuant to 40 CFR § 141.32. The system must also
report to the State pursuant to 40 CFR § 141.134. If the system fails to complete 12 consecutive months'
monitoring, compliance with the MCL for the last four-quarter compliance period must be based on an
average of the available data.
3.6 Operational Considerations
3.6.1 Process Considerations
Because ozone is such a strong oxidant, it will react with many organic and inorganic compounds
present in the water. Ozone is used to remove tastes and odors by breaking down organic
compounds, and to aid in the removal of iron and manganese by oxidizing these compounds to less
soluble forms. These demands should be satisfied before any ozone is available to satisfy primary
disinfection requirements. The presence and concentration of these compounds can dictate the
location of ozone addition, depending on the process goals.
Stolarik and Christie (1997) present the results of 10 years of operation at the 600 mgd Los Angeles
Aqueduct Filtration and Ozone Facility. Operational experiences at this facility showed lower
particle counts (greater than one micron) with ozone use. The optimum ozone concentration in the
gas phase applied was found to be 6 percent when using the cryogenic oxygen production facilities,
and 4 to 5 percent when using liquid oxygen (LOX).
3.6.2 Space Requirements
Storage of LOX is subject to regulations in building and fire codes. These regulations will impact the
space requirements and may dictate the construction materials of adjacent structures if the certain
setback requirements cannot be met. In general, the footprint for ozone generated from air is smaller
than that required for chloramination and chloride dioxide applications. However, the footprint area
for ozone generated from pure oxygen is comparable to that of chlorine dioxide because of the
additional area needed for storage.
EPA Guidance Manual 3-38 April 1999
Alternative Disinfectants and Oxidants
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3. OZONE
3.6.3 Material Selection
Ozone-resistant materials should be used from the ozone generators through the off-gas destruct unit.
If oxygen is used for the feed gas, oxygen resistant materials should be used up to the generators.
Pure oxygen piping should be specially cleaned after installation for oxygen service, which increases
construction cost. Materials for air preparation systems can be those normally used for compressed
air systems. Langlais et al. (1991) recommended that piping beyond the desiccant dryers be ozone-
resistant, as some backflow and ozone diffusion can occur. If a receiver is provided following the
desiccant dryer, the piping should be ozone-resistant, downstream of the pressure regulator. Ozone-
resistant (oxygen resistant as well if high purity oxygen is the feed gas) check valves should be
placed in the piping ahead of the generator.
Ozone-resistant materials include the austenitic (300 series) stainless steels, glass and other ceramics,
Teflon and Hypalon, and concrete. The 304 series stainless steels can be used for "dry" ozone gas
(also for oxygen), 316 series should be used for "wet" service. Wet service includes piping in the
contactors and all off-gas piping and the off-gas destruct unit. Teflon or Hypalon should be used for
gasket materials. Concrete should be manufactured from Type II or Type IV cement. Typical
practice in the United States is to provide 3 inches of cover for reinforcing to prevent corrosion by
either ozone gas or ozone in solution, although Fonlupt (1979) reports that 4 cm (1.13 inches) is
adequate for protection. Hatches for access into contactors should be fabricated from 316 series
stainless steels and provided with ozone-resistant seals.
3.6.4 Ozone System Maintenance
Stolarik and Christie (1997) provide a good overview of the operational and maintenance
requirements during the 10 years of operating the 600 mgd Los Angeles Aqueduct Filtration and
Ozone Plant. The ozone system has been available 97.1 percent of the time over the 10 year period.
Fuse failure and generator cleaning comprised the major maintenance chores on the ozone generators
during the first years. Fuse failure was caused by a malfunction when its glass dielectric tube failed.
Vessels are cleaned every three years or when exit gas temperatures rise due to Fe3O4 deposits on the
ground electrode/heat exchanger surfaces.
Rod shaped ceramic diffusers worked well as ozone diffusers for the initial two years. These were
replaced by sintered stainless steel and ultimately a modified ceramic diffuser.
3.6.5 Ozone Safety
Concern for safety even at the risk of being overcautious, would be to follow practices that have been
successfully applied to other oxidants over the years. This would be to generally isolate the
ozonation system from the remainder of the plant. This should not be interpreted to mean a separate
building, but rather separate rooms, separate exterior entrances, separate heating and ventilation
systems, noise control, etc. This method already is manifested in some of the European ozonation
plants, but on a lesser scale.
April 1999 3-39 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
3. OZONE
Ozone generators should be housed indoors for protection from the environment and to protect
personnel from leaking ozone in the case of a malfunction. Ventilation should be provided to prevent
excess temperature rise in the generator room, and to exhaust the room in the case of a leak.
Adequate space should be provided to remove the tubes from the generator shell and to service the
generator power supplies. Air prep systems tend to be noisy; therefore, it is desirable to separate
them from the ozone generators. Off-gas destruct units can be located outside if the climate is not
too extreme. If placed inside, an ambient ozone detector should be provided in the enclosure. All
rooms should be properly ventilated, heated, and cooled to match the equipment-operating
environment.
Continuous monitoring instruments should be maintained to monitor levels of ozone in the rooms.
Self-contained breathing apparatuses should be located in hallways outside the rooms liable to ozone
hazards. Ambient ozone exposure levels, which have been proposed by appropriate U.S.
organizations, are summarized below. The maximum recommended ozone levels are as follows:
• Occupational Safety and Health Administration. The maximum permissible exposure to
airborne concentrations of ozone not in excess of 0.1 mg/L (by volume) averaged over an eight-
hour work shift.
• American National Standards Institute/American Society for testing Materials
(ANSI/ASTM). Control occupational exposure such that the worker will not be exposed to
ozone concentrations in excess of a time weighted average of 0.1 mg/L (by volume) for eight
hours or more per workday, and that no worker be exposed to a ceiling concentration of ozone in
excess of 0.3 mg/L 9by volume) for more than ten minutes.
• American Conference of Government Industrial Hygienists (ACGIH). Maximum ozone
level of 0.1 mg/L (by volume) for a normal eight hour work day or 40 hour work week, and a
maximum concentration of 0.3 mg/L (by volume) for exposure of up to 15 minutes.
• American Industrial Hygiene Association. Maximum, concentration for eight hour exposure
of 0.1 mg/L (by volume).
There is a question of whether prolonged exposure to ozone may impair a worker's ability to smell or
be aware of ozone levels at less than critical levels. Awareness of an odor of ozone should not be
relied upon. Instrumentation and equipment should be provided to measure ambient ozone levels and
perform the following safety functions:
• Initiate an alarm signal at an ambient ozone level of 0.1 mg/L (by volume). Alarms should
include warning lights in the main control panel and at entrances to the ozonation facilities as
well as audible alarms.
• Initiate a second alarm signal at ambient ozone levels of 0.3 mg/L (by volume). This signal
would immediately shut down ozone generation equipment and would initiate a second set of
visual and audible alarms at the control panel and at the ozone generation facility entrances. An
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emergency ventilation system capable of exhausting the room within a period of 2 to 3 minutes
also would be interconnected to the 0.3 mg/L ozone level alarm.
Ozone gas is a hazardous gas and should be handled accordingly. Ambient ozone levels should be
monitored and equipment shut- down and alarmed when levels exceed 0.1 ppm. Emergency
ventilation is typically provided for enclosed areas. Building and fire codes will provide additional
guidance. The OSHA exposure limit for an 8-hour shift is 0.1 ppm by volume. The pungent odor of
ozone will provide warning to operators of any possible ozone leak.
3.7 Summary
3.7.1 Advantages and Disadvantages of Ozone Use
The following list highlights selected advantages and disadvantages of using ozone as a
disinfection method for drinking water (Masschelein, 1992). Because of the wide variation of
system size, water quality, and dosages applied, some of these advantages and disadvantages
may not apply to a particular system.
Advantages
• Ozone is more effective than chlorine, chloramines, and chlorine dioxide for inactivation of viruses,
Cryptosporidium, and Giardia.
• Ozone oxidizes iron, manganese, and sulfides.
• Ozone can sometimes enhance the clarification process and turbidity removal.
• Ozone controls color, taste, and odors.
• One of the most efficient chemical disinfectants, ozone requires a very short contact time.
• In the absence of bromide, halogen-substitutes DBFs are not formed.
• Upon decomposition, the only residual is dissolved oxygen.
• Biocidal activity is not influenced by pH.
Disadvantages
• DBFs are formed, particularly by bromate and bromine-substituted DBFs, in the presence of bromide,
aldehydes, ketones, etc.
• The initial cost of ozonation equipment is high.
• The generation of ozone requires high energy and should be generated on-site.
• Ozone is highly corrosive and toxic.
• Biologically activated filters are needed for removing assimilable organic carbon and biodegradable
DBFs.
• Ozone decays rapidly at high pH and warm temperatures. '
• Ozone provides no residual.
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• Ozone requires higher level of maintenance and operator skill.
3.7.2 Summary Table
Table 3-12 presents a summary of the considerations for the use of ozone as a disinfectant.
Table 3-12. Summary of Ozone Disinfection Considerations
Consideration
Generation
Primary uses
Inactivation efficiency
Byproduct formation
Limitations
Points of application
Safety considerations
Description
Because of its instability, ozone should be generated at the point of
use. Ozone can be generated from oxygen present in air or high purity
oxygen. The feed gas source should be clean and dry, with a maximum
dewpoint of -60°C. Ozone generation consumes power at a rate of 8 to
17 kWhr/kg Os. Onsite generation saves a lot of storage space.
Primary uses include primary disinfection and chemical oxidation. As an
oxidizing agent, ozone can be used to increase the biodegradability of
organic compounds destroys taste and odor control, and reduce levels
of chlorination DBP precursors. Ozone should not be used for
secondary disinfection because it is highly reactive and does not
maintain an appreciable residual level for the length of time desired in
the distribution system.
Ozone is one of the most potent and effective germicide used in water
treatment. It is effective against bacteria, viruses, and protozoan cysts.
Inactivation efficiency for bacteria and viruses is not affected by pH; at
pH levels between 6 and 9. As water temperature increases, ozone
disinfection efficiency increases.
Ozone itself does not form halogenated DBPs; however, if bromide ion
is present in the raw water or if chlorine is added as a secondary
disinfectant, halogenated DBPs, including bromate ion may be formed.
Other ozonation byproducts include organic acids and aldehydes.
Ozone generation is a relatively complex process. Storage of LOX (if
oxygen is to be the feed gas) is subject to building and fire codes.
For primary disinfection, ozone addition should be prior to
biofiltration/filtration and after sedimentation. For oxidation, ozone
addition can be prior to coagulation/sedimentation or filtration
depending on the constituents to be oxidized.
Ozone is a toxic gas and the ozone production and application facilities
should be designed to generate, apply, and control this gas, so as to
protect plant personnel. Ambient ozone levels in plant facilities should
be monitored continuously.
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3.8 References
1. Alceon Corp. 1993. Overview of Available Information on the Toxicity of Drinking Water
Disinfectants and Their By-products. Cambridge, MA.
2. Amy, G.L., M.S. Siddiqui. 1991. "Ozone-Bromide Interactions in Water Treatment."
Conference proceedings, AWWA Annual Conference, Philadelphia, PA.
3. AWWA (American Water Works Association). 1990. Water Quality and Treatment. F.W.
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Rate Biological Filtration." J. AWWA.80(12):47.
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Water Treatment Application and Engineering. AWWARF.
6. Billen, G., et al. 1985. Action des Populations Bacteriennes Vis-a-Vis des Matieres Organiques
dans les Filtres Biologiques. Report to Compagnie Generate des Eaux, Paris.
7. Boyce, D.S., et al. 1981. "The Effect of Bentonite Clay on Ozone Disinfection of Bacteria and
Viruses in Water." Water Res. 15:759-767.
8. Bringmann, G. 1954. "Determination of the Lethal Activity of Chlorine and Ozone on E. coli."
Z. f., Hygiene. 139:130-139.
9. Chang, S.L. 1971. "Modern Concept of Disinfection." J. Sanit. Engin. Division. 97:689-707.
10. Cronholm, L.S., et all 976. "Enteric Virus Survival in Package Plants and the Upgrading of the
Small Treatment Plants Using Ozone." Research Report No. 98, Water Resources Research
Institute, University of Kentucky, Lexington, KY.
11. DeMers, L.D. and R.C. Renner. 1992. Alternative Disinfection Technologies for Small Drinking
Water Systems. AWWARF and AWWA, Denver, CO.
12. Dimitriou, M.A. (editor). 1990. Design Guidance Manual for Ozone Systems. International
Ozone Association, Norwalk, CN.
13. Domingue, E. L., et al. 1988. "Effects of Three Oxidizing Biocides on Legionella pneumophila,
Serogroup 1." Appl. Environ. Microbiol. 40:11-30.
14. Eighmy, T.T., S.K. Spanos, J. Royce, M.R. Collins, J.P. Malley. 1991. "Microbial Activity in
Slow Sand Filters." Conference proceedings, Slow Sand Filtration Workshop, Timeless
Technology for Modern Applications, Durham, NH.
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15. Farooq, S. et al.1977. "The Effect of Ozone Bubbles on Disinfection." Progr. Water Ozone Sci.
E?ig. 9(2):233.
16. Farooq, S. 1976. Kinetics of Inactivation of Yeasts and Acid-Fast Organisms with Ozone. Ph.D.
Thesis, University of Illinois at Urbana-Champaign, IL.
17. Farvardin, M.R. and A.G. Collins. 1990. "Mechanism(s) of Ozone Induced Coagulation of
Organic Colloids." Conference proceedings, AWWA Annual Conference, Cincinnati, OH.
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Francisco, CA.
19. Franson, M.H., Eaton, A.D., Clesceri, L.S., and Greenberg, A.E., (editors). 1995. Standard
Methods for the Examination of Water and Wastewater, Nineteenth Edition. American Public
Health Association, AWWA, and Water Environment Federation, Washington D.C.
20. Georgeson, D.L. and A.A. Karimi. 1988. "Water Quality Improvements with the Use of Ozone
at the Los Angeles Water Treatment Plant." Ozone Sci. Engrg. 10(3):255-276.
21. Giese, A.C. and E. Christensen. 1954. "Effects of Ozone on Organisms." Physiol. Zool. 27:101.
22. Glaze W.H., M. Koga, D. Cancilla. 1989a. "Ozonation Byproducts. 2. Improvement of an Aqueous-
Phase Derivatizition Method for the Detection of Formaldehyde and Other Carbonyl Compounds
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25. Glaze, W.H., and J.W. Kang. 1988. "Advanced Oxidation Processes for Treating Groundwater
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26. Glaze, W.H., H.S. Weinberg, S.W. Krasner, M.J. Sclimenti. 1991. "Trends in Aldehyde Formation
and Removal Through Plants Using Ozonation and Biological Active Filters." Conference
proceedings, AWWA, Philadelphia, PA.
27. Goldstein, B.D., and E.M. McDonagh. 1975. "Effect of Ozone on Cell Membrane Protein
Fluorescence I. in vitro Studies Utilizing the Red Cell Membrane." Environ. Res. 9:179-186.
28. Gordon, G. K. Rankness, D. Vornehm, and D. Wood. 1989. "Limitations of the lodometric
Determination of Ozone." J. AWWA. 81(6):72-76.
29. Gordon, G., W.J. Cooper, R.G. Rice, and G.E. Pacey. 1992. Disinfectant Residual Measurement
Methods, second edition. AWWARF and AWWA, Denver, CO.
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30. Gurol, M.D. and M. Pidatella. 1983. "A Study of Ozone-Induced Coagulation." Conference
proceedings, ASCE Environmental Engineering Division Specialty Conference. Allen Medine
and Michael Anderson (editors), Boulder, CO.
31. Haag, W.R. and J. Hoigne. 1984. "Kinetics and Products of the Reactions of Ozone with
Various Forms of Chlorine and Bromine in Water." Ozone Sci. Engrg. 6(2): 103-14.
32. Hann, V.A. 1956. "Disinfection of Drinking Water with Ozone." J. AWWA. 48(10): 1316.
33. Harakeh, M.S. and M. Butler. 1984. "Factors Influencing the Ozone Inactivation of Enteric
Viruses in Effluent." Ozone Sci. Engrg. 6:235-243.
34. Hiltebrand, D.J., A.F. Hess, P.B. Galant, and C.R. O'Melia. 1986. "Impact of Chlorine Dioxide
and Ozone Preoxidation on Conventional Treatment and Direct Filtration Treatment Processes."
Conference proceedings, AWWA Seminar on Ozonation: Recent Advances and Research
Needs, Denver, CO.
35r Hoff, J.C. 1986. Inactivation of Microbial Agents by Chemical Disinfectants, U. S.
Environmental Protection Agency, EPA/600/2-86/067.
36. Hoigne J. and H. Bader. 1976. Role of Hydroxyl Radical Reactions in Ozonation Processes in
Aqueous Solutions, Water Res. 10: 377.
37. Hoigne J., and H. Bader. 1988. "The Formation of Trichloronitromethane (chloropicrin) and
Chloroform in a Combined Ozonation/Chlorination Treatment of Drinking Water." Water Res.
22(3):313.
38. Hoigne J., and H. Bader. 1983b. "Rate Constants of Reaction of Ozone with Organic and
Inorganic Compounds in Water - II. Dissociating Organic Compounds." Water Res. 17:185-194.
39. Hoigne J., and H. Bader. 1977. "The Role of Hydroxyl Radical Reactions in Ozonation
Processes in Aqueous Solutions." Water Res. 10:377-386.
40. Hoigne J., and H. Bader. 1983a. "Rate Constants of Reaction of Ozone with Organic and
Inorganic Compounds in Water -1. Non-dissociating Organic Compounds." Water Res. 17:173-
183.
41. Hoigne, J. and Bader, H. 1975. Ozonation of Water: Role of Hydroxyl Radicals as Oxidizing
Intermediates. Science, Vol. 190, pp. 782.
42. Huck, P.M., P.M. Fedorak, and W.B. Anderson. 1991. "Formation and Removal of Assimilable
Organic Carbon During Biological Treatment." /. AWWA. 83(12):69-80.
43. IOA (International Ozone Association). 1989. Photometric Measurement of Low Ozone
Concentrations in the Gas Phase. Standardisation Committee-Europe.
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44. Katz, J. 1980. Ozone and Chlorine Dioxide Technology for Disinfection of Drinking Water.
Noyes Data Corporation, Park Ridge, New Jersey.
45. Katzenelson, E., et al. 1974. "Inactivation Kinetics of Viruses and Bacteria in Water by Use of
Ozone." J. AWWA. 66:725-729.
46. Keller, J.W., R.A. Morin, and T.J. Schaffernoth. 1974. "Ozone Disinfection Pilot Plants Studies
at Laconia, New Hampshire." J. AWWA. 66:730.
47. Kim, C.K., et al. 1980. "Mechanism of Ozone Inactivation of Bacteriophage f2." Appl. Environ.
Microbiol 39:210-218.
48. Kinman, R.N. 1975. "Water and Wastewater Disinfection with Ozone: A Critical Review." Crit.
Rev. Environ. Contr. 5:141-152.
49. Krasner, S.W., W.H. Glaze, H.S. Weinberg, et al. 1993. "Formation of Control of Bromate
During Ozonation of Water Containing Bromide." J. AWWA. 85(5):62..
50. Krasner, S.W., et al. 1989. "The Occurrence of Disinfection By-products in US Drinking
Water." J. 4WVK4.81(8):41.
51. Langlais, B., et al. 1990. "New Developments: Ozone in Water and Wastewater Treatment. The
CT Value Concept for Evaluation of Disinfection Process Efficiency; Particular Case of
Ozonation for Inactivation of Some Protozoa, Free-Living Amoeba and Cryptosporidium."
Presented at the Int. Ozone Assn. Pan-American Conference, Shreveport, Louisiana, March 27-
29.
52. Langlais, B., D.A. Reckhow, and D.R. Brink (editors). 1991. Ozone in Drinking Water
Treatment: Application and Engineering. AWWARF and Lewis Publishers, Boca Raton, FL.
53. Langlais B. and D. Perrine. 1986. "Action of Ozone on Trophozoites and Free Amoeba Cysts,
Whether Pathogenic of Not." Ozone Sci. Engrg. 8(3): 187-198.
54. LeChevallier, M.W., W.C. Becker, P. Schorr, and R.G. Lee. 1992. "Evaluating the Performance of
Biologically Active Rapid Filters." J. AWWA. 84(4): 136-146.
55. LeLacheur, R.M., P.C. Singer, and M J. Charles. 1991. "Disinfection By-products in New Jersey
Drinking Waters." Conference proceedings, AWWA Annual Conference, Philadelphia, PA.
56. Logsdon, G.S., S. Foellmi, and B. Long. 1992. "Filtration Pilot Plant Studies for Greater
Vancouver's Water Supply." Conference proceedings, AWWA Annual Conference, Vancouver,
British Columbia.
57. Malley, J.P., T.T. Eighmy, M.R. Collins, J.A. Royce, and D.F. Morgan. 1993. "The True
Performance and Microbiology of Ozone - Enhanced Biological Filtration." J. AWWA.
85(12):47-57.
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58. Masschelein, W.J. 1992. "Unit Processes in Drinking Water Treatment." Marcel Decker D.C.,
New York , Brussels, Hong Kong.
59. Masschelein, W.J. 1977. "Spectrophotometric Determination of Residual Ozone in Water with
ACVK." J. AWWA. 69:461-462.
60. McGuire, M.J., S.W. Krasner, and J. Gramith. 1990. Comments on Bromide Levels in State
Project Water and Impacts on Control of Disinfection Byproducts Metropolitan Water District
of Southern California.
61. McKee, H.C., R.E. Childers, and V.B. Parr 1975. Collaborative Study of Reference Method for
Measurement of Photochemical Oxidants in the Atmosphere, EPA EPA-650/4-75-016,
Washington, D.C. February.
62. McKnight A., and D.A. Reckhow. 1992. "Reactions of Ozonation Byproducts with Chlorine and
Chloramines." Conference proceedings, AWWA Annual Conference, Vancouver, British
Columbia.
63. MWDSC and JMM (Metropolitan Water District of Southern California and James M.
Montgomery Consulting Engineers). 1992. "Pilot Scale Evaluation of Ozone and peroxone."
AWWARF and AWWA, Denver, CO.
64. Morris, J.C. 1975. "Aspects of the Quantitative Assessment of Germicidal Efficiency."
Disinfection: Water and Wastewater. J.D. Johnson (editor). Ann Arbor Science Publishers, Inc.,
Ann Arbor, MI.
65. Owens, J. H., et al. 1994. "Pilot-Scale Ozone Inactivation of Cryptosporidium and Giardia"
Conference proceedings, Water Quality Technology Conference, Part II, San Francisco, CA.
66. Peeters, J. E. et al. 1989. "Effect of Disinfection of Drinking Water with Ozone or Chlorine
Dioxide on Survival of Cryptosporidium parvum Oocysts." Appl. Environ. Microbiol.
55(6): 1519-1522.
67. Perrine, D., et al. 1984. "Action d 1'Ozone sur les Trophozoites d'Amibes Libres Pathogens ou
Non." Bull Soc.Frnac. Parasitol. 3:81.
68. Prendiville, D.A. 1986. "Ozonation at the 900 cfs Los Angeles Water Purification Plant." Ozone
Sci. Engrg. 8:77.
69. Price, M.L. 1994. Ozone and Biological Treatment for DBF Control and Biological Stability.
AWWARF and AWWA, Denver, CO, pp. 252.
70. Rachwal, A.J., et al. 1988. "Advanced Techniques for Upgrading Large Scale Slow Sand
Filters." Slow Sand Filtration- Recent Developments in Water Treatment Technology, Ellis
Horwood Ltd, Chichester, U.K.
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71. Reckhow, D.A., J.K. Edzwald, and I.E. Tobiason. 1993. "Ozone as an Aid to Coagulation and
Filtration." AWWARF and AWWA, Denver, CO.
72. Reckhow, D.A., P.C.Singer, and R.R. Trussell. 1986. Ozone as a coagulant aid. Seminar
proceedings, Ozonation, Recent Advances and Research Needs, AWWA Annual Conference,
Denver, CO.
73. Reckhow, D.A., I.E. Tobiason, M.S. Switzenbaum, R. McEnroe, Y. Xie, X. Zhou, P.
McLaughlin, and H.J. Dunn. 1992. "Control of Disinfection Byproducts and AOC by Pre-
Ozonation and Biologically Active In-Line Direct Filtration." Conference proceedings, AWWA
Annual Conference, Vancouver, British Columbia.
74. Regli, S., J.E. Comwell, X. Zhang., et al. 1992. Framework for Decision Making: An EPA
Perspective. EPA 81 l-R-92-005, EPA, Washington, D.C.
75. Rehme, K.A., J.C. Puzak, M.E. Beard, C.F. Smith, and R.J. Paur. 1980. Evaluation of Ozone
Calibration Procedures, EPA-600/S4-80-050, EPA, Washington, D.C, February.
76. Renner, R.C., M.C. Robson, G.W. Miller, and A.G. Hill. 1988. "Ozone in Water Treatment -
The Designer's Role." Ozone Sci. Engrg. 10(l):55-87.
77. Rice, R.G. 1996. Ozone Reference Guide. Electric Power Research Institute, St. Louis, MO.
78. Rice, R.G., P.K. Overbeck, K. Larson. 1998. Ozone Treatment for Small Water Systems.
Presented at the First International Symposium on Safe Drinking water in Small Systems. NSF
International/PAHPAVHO, Arlington, VA, May 10-13, 1998.
79. Riesser, V.W., et al. 1976. "Possible Mechanisms of Poliovirus Inactivation by Ozone." Forum
on Ozone Disinfection, E. G. Fochtman, R.G. Rice, and M.E. Browning (editors), pp. 186-192,
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80. Rittman, B.E. 1990. "Analyzing Biofilm Processes Used in Biological Filtration." J. AWWA.
82(12):62.
81. Roy, D. 1979. Inactivation of Enteroviruses by Ozone. Ph.D. Thesis, University of Illinois at
Urbana-Champaign.
82. Roy, D., R.S. Engelbrecht, and E.S.K. Chian. 1982. "Comparative Inactivation of Six
Enteroviruses by Ozone." J. AWWA. 74(12):660.
83. Scott, D.B.M. and E.G. Lesher. 1963. "Effect of Ozone on Survival and Permeability of
Escherichia coli." J. Bacterial. 85:567-576.
84. Singer P.C. 1992. "Formation and Characterization of Disinfection Byproducts." Presented at
the First International Conference on the Safety of Water Disinfection: Balancing Chemical and
Microbial Risks.
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85. Singer, P.C., et al. 1989. "Ozonation at Belle Glade, Florida: A Case History." Conference
proceeding, IOA Ninth Ozone World Conference.
86. Song, R., et al. 1997. "Bromate Minimization During Ozonation." J. AWWA. 89(6):69.
87. Sproul, O. J., et al. 1982. "The Mechanism of Ozone Inactivation of Waterborne Viruses."
Water Sci. Technol. 14:303-314.
88. Staehelin, J., R.E. Biihler, and J. Hoigne. 1984. "Ozone Decomposition in Water Studies by
Pulse Radiolysis. 2 OH and HO4 as Chain Intermediates." J. Phys. Chem. 88:5999-6004.
89. Stolarik, G. F., and J.D. Christie. 1997. "A Decade of Ozonation in Los Angeles." Conference
proceedings, IOA Pan American Group Conference, Lake Tahoe, NV.
90. Suffet, I.H., C. Anselme, and J. Mallevialle. 1986. "Removal of Tastes and Odors by
Ozonation." Conference proceedings, AWWA Seminar on Ozonation: Recent Advances and
Research Needs, Denver, CO.
91. Tobiason, J.E., J.K. Edzwald, O.D. Schneider, M.B. Fox, and H.J. Dunn. 1992. "Pilot Study of
the Effects of Ozone and Peroxone on In-Line Direct Filtration." /. AWWA. 84(12):72-84.
92. Tomiyasu, H., and G. Gordon. 1984. "Colorimetric Determination of Ozone in Water Based on
Reaction with Bis-(terpyridine)iron(II)." Analytical Chem. 56:752-754.
93. Troyan, J.J. and S.P. Hansen. 1989. Treatment ofMicrobial Contaminants in Potable Water
Supplies Technologies and Costs. Noyes Data Corporation, Park Ridge, New Jersey.
94. Umphries, M.D., et al. 1979. "The Effects of Pre-ozonation on the Formation of
Trihalomethanes." Ozonews. 6(3).
95. USEPA. 1997. USEPA Method 300.1, Determination of Inorganic Anions in Drinking Water
by Ion Chromatography, Revision 1.0. EPA A/600/r-98/l 18.
96. Van Dijk, J.F.M., and R.A. Falkenberg. 1977. "The Determination of Ozone Using the Reaction
with Rhodamine B/Gallic Acid." Presented at Third Ozone World Congress sponsored by the
IOA, Paris, France, May.
97. Van Gunten, U. and J. Hoigne. 1996. "Ozonation of Bromide-Containing Waters: Bromate
Formation through Ozone and Hydroxyl Radicals." Disinfection By-Products in Water
Treatment, Minear, R.A. and G.L. Amy (editors). CRC Press, Inc., Boca Raton, FL.
98. Van Hoof, F., J.G. Janssens, H. van Dijck. 1985. "Formation of Mutagenic Activity During
Surface Water Pre-ozonation and Its Removal in Drinking Water Treatment." Chemosphere,
14(5):501.
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99. Vaughn, J.M., et al. 1987. "Inactivation of Human and Simian Rotaviruses by Ozone." Appl.
Env. Microbiol. 53:2218-2221.
100. Walsh, D.S., et al. 1980. "Ozone Inactivation of Floe Associated Viruses and Bacteria." J.
Environ. Eng. Div. ASCE. 106:711-726.
101. Ward, S.B. and D.W. Larder. 1973. "The Determination of Ozone in the Presence of Chlorine."
Water Treatment Examination. 22:222-229.
102. Wickramanayake, G.B., et al. 1984b. "Inactivation of Giardia lamblia Cysts with Ozone." Appl.
Env. Microbiol. 48(3):671-672.
103. Wickramanayake, G.B., et al. 1984a. "Inactivation of Naegleria and Giardia cysts in Water by
Ozonation." J. Water Pollution Control Fed. 56(8):983-988.
104. Wickramanayake, G.B. 1984c. Kinetics and Mechanism of Ozone Inactivation of Protozoan
Cysts. Ph.D. dissertation, Ohio State University, Columbus, OH.
105. Wuhrmann, K., and J. Meyrath. 1955. "The Bactericidal Action of Ozone Solution. Schweitz."
J. Allgen. Pathol. Bakteriol. ,18:1060.
106. Yamada, H. and I. Somiya. 1989. "The Determination of Carbonyl Compounds in Ozonated
Water By the PFBOA Method." Ozone Sci. Engrg. 11 (2): 127.
107. Zabel, T.F. 1985. "The Application of Ozone for Water Treatment in the United Kingdom -
Current Practice and Recent Research." Ozone Sci. Engrg. 7(1):11.
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Since the beginning of the twentieth century, when it was first used at a spa in Ostend, Belgium,
chlorine dioxide has been known as a powerful disinfectant of water. During the 1950s, it was
introduced more generally as a drinking water disinfectant since it provided less organoleptic
hindering than chlorine. Approximately 700 to 900 public water systems use chlorine dioxide to
treat potable water (Hoehn, 1992). Today, the major uses of chlorine dioxide are:
• CT disinfection credit;
• Preoxidant to control tastes and odor;
• Control of iron and manganese; and
• Control of hydrogen sulfide and phenolic compounds.
4.1 Chlorine Dioxide Chemistry
4.1.1 Oxidation Potential
The metabolism of microorganisms and consequently their ability to survive and propagate are
influenced by the oxidation reduction potential (ORP) of the medium in which it lives (USEPA,
1996).
Chlorine dioxide (C1O2) is a neutral compound of chlorine in the +IV oxidation state. It disinfects by
oxidation; however, it does not chlorinate. It is a relatively small, volatile, and highly energetic
molecule, and a free radical even while in dilute aqueous solutions. At high concentrations, it reacts
violently with reducing agents. However, it is stable in dilute solution in a closed container in the
absence of light (AWWA, 1990). Chlorine dioxide functions as a highly selective oxidant due to its
unique, one-electron transfer mechanism where it is reduced to chlorite (C1O2") (Hoehn et al., 1996).
The pKa for the chlorite ion, chlorous acid equilibrium, is extremely low at pH 1.8. This is
remarkably different from the hypochlorous acid/hypochlorite base ion pair equilibrium found near
neutrality, and indicates the chlorite ion will exist as the dominant species in drinking water. The
oxidation reduction of some key reactions are (CRC, 1990):
C/02(aq) + e = ClOi E° = 0.954V
Other important half reactions are:
E° = 0.76V
E° = 0.33V
ClOi + 2H+ + e = C102 + H20 E° = 1.152V
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4. CHLORINE DIOXIDE
In drinking water, chlorite (ClOi") is the predominant reaction endproduct, with approximately 50 to
70 percent of the chlorine dioxide converted to chlorite and 30 percent to chlorate (CICV) and
chloride (Cl") (Werdehoff and Singer, 1987).
4.2 Generation
4.2.1 Introduction
One of the most important physical properties of chlorine dioxide is its high solubility in water,
particularly in chilled water. In contrast to the hydrolysis of chlorine gas in water, chlorine dioxide
in water does not hydrolyze to any appreciable extent but remains in solution as a dissolved gas
(Aieta and Berg, 1986). It is approximately 10 times more soluble than chlorine (above 11°C), while
it is extremely volatile and can be easily removed from dilute aqueous solutions with minimal
aeration or recarbonation with carbon dioxide (e.g. softening plants). Above 11 to 12°C, the free
radical is found in gaseous form. This characteristic may affect chlorine dioxide's effectiveness when
batching solutions and plumbing appropriate injection points. Other concerns are the increased
difficulty in analyzing for specific compounds in the presence of many interfering
compounds/residual longevity and volatility of gaseous compounds. In the gaseous form, the free
radicals also react slowly with water. The reaction rate is 7 to 10 million times slower than that of
the hydrolysis rate for chlorine gas (Gates, 1989).
Chlorine dioxide cannot be compressed or stored commercially as a gas because it is explosive under
pressure. Therefore, it is never shipped. Chlorine dioxide is considered explosive at higher
concentrations which exceed 10 percent by volume in air, and its ignition temperature is about 130°C
(266°F) at partial pressures (National Safety Council Data Sheet 525 - C1O2, 1967). Strong aqueous
solutions of chlorine dioxide will release gaseous chlorine dioxide into a closed atmosphere above
the solution at levels that may exceed critical concentrations. Some newer generators produce a
continuous supply of dilute gaseous chlorine dioxide in the range of 100 to 300 mm-Hg (abs) rather
than in an aqueous solution (National Safety Council, 1997). For potable water treatment processes,
aqueous solutions between 0.1 and 0.5 percent are common from a number of current generation
technologies.
Most commercial generators use sodium chlorite (NaClO2) as the common precursor feedstock
chemical to generate chlorine dioxide for drinking water application. Recently, production of
chlorine dioxide from sodium chlorate (NaClO3) has been introduced as a generation method where
in NaClOa is reduced by a mixture of concentrated hydrogen peroxide (H2O2) and concentrated
sulfuric acid (HiSO.)). Chlorate-based systems have traditionally been used in pulp and paper
applications, but have recently been tested full-scale at two U.S. municipal water treatment plants.
This is an emerging technology in the drinking water field and is not discussed in this guidance
manual.
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4. CHLORINE DIOXIDE
4.2.2 Chlorine Dioxide Purity
Chlorine dioxide generators are operated to obtain the maximum production (yield) of chlorine
dioxide, while minimizing free chlorine or other residual oxidant formation. The specified yield for
chlorine dioxide generators is typically greater than 95 percent. In addition, the measurable excess
chlorine should be less than 2 percent by weight in the generator effluent. Generator yield is defined
as (Gordon et al., 1990):
Yield = f n . L V ,,»^r ix 100
[cio2]+[cio-2\+(i£)[ao;\
Where: [ CIO2 ] = Chlorine dioxide concentration, mg/L.
[ ClO^ ] = Chlorite concentration, mg/L.
[ ClO^ ] = Chlorate concentration, mg/L.
' 67.45N
83.45
= Molecular weight ratio of C1O2" to C1O3".
Since any chlorite ion fed to the generator may result in the formation of C1O2, C1O2", or C1O3~, the
purity of the resultant mixture can be calculated using the concentrations of each of the species from
appropriate analytical measurements. The determination of purity requires neither flow
measurement, mass recoveries, nor manufacturer-based methods to determine production "yield,"
"theoretical yield," "efficiency," or conversion for any precursor feedstock. This approach does not
require flow measurements that can introduce up to 5 percent error in the calculations.
Utilities that use chlorine dioxide should measure excess chlorine (as FAC) in the generator effluent
in addition to the C1O2" related species. FAC may appear.as false C1O2 residuals for CT purposes, or
result in the formation of chlorinated DBFs if high, relative to the C1O2 level in the generated
mixture. Excess chlorine is defined as:
Excess Cl 2 =
[C19]
[C102] + [C10-] + [C103 ] x
Where: - '- - = stoichiometric and molecular weight ratio of C12 to C1O2".
(2x67.45)
The following represents a summarily simpler equation that substantially resolves the problems of
different equipment-specific calibration methods, chlorine-contaminated C1O2, or low efficiency
conversion of either chlorite- or chlorate-based precursor material.
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4. CHLORINE DIOXIDE
[C109]
Purity = ^ x 100
[C1O2 ] + [FAC] + [C10~ ] + [C10~ ]
This practical (weight-based) calculation permits a variety of approved analytical methods (discussed
in section 4.6) to be used to assess generator performance on unbiased scientific principles, rather
than non-standardized manufacturer specifications.
4.2.3 Methods of Generating Chlorine Dioxide
For potable water application, chlorine dioxide is generated from sodium chlorite solutions. The
principal generation reactions that occur in the majority of generators have been known for a long
time. Chlorine dioxide can be formed by sodium chlorite reacting with gaseous chlorine (Cl2(g)),
hypochlorous acid (HOC1), or hydrochloric acid (HC1). The reactions are:
2NaClO2 + Cl2(g) = 2C102(g) + 2NaCl [ 1 a]
2NaClO2 + HOC1 = 2ClO2(g) + NaCl + NaOH [ 1 b]
5NaClO2 -f 4HC1 = 4ClO2(g) + SNaCl + 2H2O [ 1 c]
Reactions [la], [Ib], and [Ic] explain how generators can differ even though the same feedstock
chemicals are used, and why some should be pH controlled and others are not so dependent on low
pH. In most commercial generators, there may be more than one reaction taking place. For example,
the formation and action of hypochlorous acid as an intermediate (formed in aqueous solutions of
chlorine) often obscures the "overall" reaction for chlorine dioxide production.
Table 4-1 provides information on some types of available commercial generators. Conventional
systems react sodium chlorite with either acid, aqueous chlorine, or gaseous chlorine. Emergent
technologies identified in Table 4-1 include electrochemical systems, a solid chlorite inert matrix
(flow-through gaseous chlorine) and a chlorate-based emerging technology that uses concentrated
hydrogen peroxide and sulfuric acid.
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4. CHLORINE DIOXJDE
Table 4-1. Commercial Chlorine Dioxide Generators
GENERATOR TYPE
MAIN REACTIONS
Reactants, byproducts, key
reactions, and chemistry notes
SPECIAL ATTRIBUTES
ACID-CHLORITE:
(Direct Acid System)
4HCI + 5NaCI02 •* 4CI02(aq) + ClOr
• Low pH
• ClOa" possible
• Slow reaction rates
Chemical feed pump interlocks required.
Production limit ~ 25-30 Ib/day.
Maximum yield at -80% efficiency.
AQUEOUS CHLORINE-
CHLORITE:
(Cl2 gas ejectors with chemical
pumps for liquids or booster
. pump for ejector water).
H20-»[HOCI/HCI]
[HOCI/HCI] + NaCI02 •»
CI02(g> + H/OCI- + NaOH + ClOr
« LowpH
• CIOs' possible
• Relatively slow reaction rates
Excess CI2 or acid to neutralize NaOH.
Production rates limited to ~ 1000 Ib/day.
High conversion but yield only 80-92%
More corrosive effluent due to low pH (-2.8-3.5).
Three chemical systems pump HC1,
hypochlorite, chlorite, and dilution water to
reaction chamber.
RECYCLED AQUEOUS
CHLORINE OR "FRENCH
LOOP"™
(Saturated CI2 solution via a
recycling loop prior to mixing
with chlorite solution.)
2HOCI + 2NaCl02
2NaOH
2CI02 + CI2
• Excess CI2 or HCI needed due to
NaOH formed.
Concentration of -3 g/L required for maximum
efficiency.
Production rate limited to - 1000 Ib/day.
Yield of 92-98% with ~10% excess CI2 reported.
Highly corrosive to pumps; draw-down
calibration needed. Maturation tank required
after mixing.
GASEOUS CHLORINE-
CHLORITE
(Gaseous CI2 and 25% solution
of sodium chlorite; pulled by
ejector into the reaction column.]
CI2(g) + NaCI02(aq) •» CI02(aq)
• Neutral pH
• Rapid reaction
• Potential scaling in reactor under
vacuum due to hardness of
feedstock.
Production rates 5-120,000 Ib/day.
Ejector-based, with no pumps. Motive water is
dilution water. Near neutral pH effluent. No
excess CI2. Turndown rated at 5-1 OX with yield
of 95-99%. Less than 2% excess CI2. Highly
calibrated flow meters with min. line pressure -
40 psig needed.
GASEOUS CHLORINE-
SOLIDS CHLORITE MATRIX
(Humidified CI2 gas is pulled or
pumped through a stable matrix
containing solid sodium chlorite.)
CI2(g> + NaCI02(S) •» CI02(g) + NaCI
• Rapid reaction rate
• New technology
Cl2 gas diluted with N2 or filtered air to produce
-8% gaseous CI02 stream. Infinite turndown is
possible with >99% yield. Maximum rate to
-1200 Ib/day per column; ganged to >10,000
Ib/day.
ELECTROCHEMICAL
(Continuous generation of Cl02
from 25% chlorite solution
recycled through electrolyte cell)
NaCI02(aq) •> CI02(aq) + 6"
• New technology
Counter-current chilled water stream accepts
gaseous CI02 from production cell after it
diffuses across the gas permeable membrane.
Small one-pass system requires precise flow for
power requirements (Coulombs law).
ACID/PEROXIDE/CHLORIDE
2NaClOs + H202 + H2S04
02 + NaS04 + H20
• 2CI02 + Uses concentrated H202 and H2S04.
Downscaled version; Foam binding; Low pH.
Source: Adapted from Gates, 1998.
4.2.3.1 Commercial Generators
The conventional chlorine-chlorite solution method generates chlorine dioxide in a two-step process.
First, chlorine gas is reacted with water to form hypochlorous acid and hydrochloric acid. These
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4. CHLORINE DIOXIDE
acids then react with sodium chlorite to form chlorine dioxide. The ratio of sodium chlorite to
hypochlorous acid should be carefully controlled. Insufficient chlorine feed will result in a large
amount of unreacted chlorite. Excess chlorine feed may result in the formation of chlorate ion,
which is an oxidation product of chlorine dioxide and not currently regulated.
Acid-Chlorite Solution - Chlorine dioxide can be generated in direct-acidification generators by
acidification of sodium chlorite solution. Several stoichiometric reactions have been reported for
such processes (Gordon et al., 1972). When chlorine dioxide is generated in this way, hydrochloric
acid is generally preferred (Reaction [lc]).
Aqueous Chlorine-Chlorite Solution - Chlorite ion (from dissolved sodium chlorite) will react with
hydrochloric acid and hypochlorous acid to form chlorine dioxide in these systems, commonly
referred to as conventional systems (Reaction [lb]):
Figure 4-1 shows a typical chlorine dioxide generator using aqueous chlorine-chlorite solution
(Demers and Renner, 1992).
If chlorine gas and chlorite ion are allowed to react under ideal conditions (not usually formed in
aqueous chlorine type systems), the resulting pH of the effluent may be close to 7. To fully utilize
sodium chlorite solution, the more expensive of the two ingredients, excess chlorine is often used.
This approach lowers the pH and drives the reaction further toward completion. The reaction is
faster than the acid-chlorite solution method, but much slower than the other commercial methods
described in the following discussion.
Recycled Aqueous Chlorine or "French Loop"™ - In this aqueous chlorine design, shown in
Figure 4-2, chlorine gas is injected into a continuously circulating water loop. This eliminates the
need for a great excess of Cb gas to be fed to the generator since the molecular chlorine will dissolve
in the feed water, and thus maintain a low pH level of the feed water. Loop-based generators keep
chlorine at or above saturation levels. The low pH condition results in high yields of chlorine
dioxide (greater than 95 percent at design production rate) (Thompson, 1989). Chlorine in the
generator effluent may react with chlorine dioxide to form chlorate if allowed to stand in batch
storage too long. The "French Loop" type of generator is more difficult to operate due to system
start-up and control of sodium chlorite feed rate (meter pump), chlorine feed rate (rotameter), and the
recirculating loop (pump). Newer designs incorporate a second batching tank for continuous
aqueous chlorine storage, thus removing many of these startup or recycling difficulties.
Gaseous Chlorine-Chlorite Solution - Sodium chlorite solution can be "vaporized" and reacted
under vacuum with molecular gaseous chlorine. This process uses undiluted reactants and is much
more rapid than chlorine solutionxhlorite solution methods (Pitochelli, 1995). Production rates are
more easily scaled up, and some installed systems have reported producing more than 60,000 pounds
per day.
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4. CHLORINE DIOXIDE
Control Box
Chlorine Dioxide
Solution
Sight Glass
Vacuum I°P11 ^T
:ulator .H^ ^Ven<
Reg
with Loss of
Chlorine Switch
Chlorine
Gas
'
water Inlet
Source: Demersand Renner, 1992.
Figure 4-1. Conventional Chlorine Dioxide Generator When Using
Chlorine-Chlorite Method
The acid-sodium hypochlorite-sodium chlorite method of generating chlorine dioxide is used when
chlorine gas is not available. First, sodium hypochlorite is combined with hydrochloric or another
acid to form hypochlorous acid. Sodium chlorite is then added to this reaction mixture to produce
chlorine dioxide.
4.2.3.2 pH Effects on Chlorine Dioxide Generation
If hypochlorous acid is formed, one of the byproducts of its reaction with sodium chlorite in solution
is sodium hydroxide. Since sodium hydroxide is also a common stabilizer of sodium chlorite
feedstock, the resulting pH of the mixture can be too high. A high pH slows the formation of
chlorine dioxide and impels less efficient chlorate-forming reactions. This is the same process in
which chlorite and hypochlorite ions react in drinking water to form chlorate ion. This neutralizing
effect of caustic may be influenced by different stabilities used in each of the types and sources of
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4. CHLORINE DIOXIDE
sodium chlorite which are approved for use in drinking water under AWWA Standard B303-95
(AWWA, 1995).
QtfeMinc Dioxide
SukKKut
Sigh* Glass
Control Box
Chlorine Dioxide Wall Cabinet | !
-feh
-Level
Switch
Water
Flowmcier
Pressure
Regulator
Bypass
Line
109 Active
Sodium Chltxilc
SolutkHi Tank
Source: Darners and Renner, 1992.
Figure 4-2. Chlorine Dioxide Generation Using Recycled Aqueous Chlorine
Method
In very low pH aqueous chlorine solutions, chlorous acid (and not the chlorite ion) may be directly
oxidized to chlorine dioxide as shown in reaction [Id]. At this low pH, gaseous chlorine remains
"dissolved" in the water at concentrations higher than the normal occurrence, and allows reaction
[la] to proceed.
4.2.3.3 Chlorate Byproduct Formation
One of the most undesirable byproducts in generators is the chlorate ion (C1O3"). Chlorate
production is possible through reactions with the intermediate dimer, {C12O2}. Rather than the
chlorite ion being simply "converted" to chlorine dioxide, reactions [la] through [Id] can result in the
supposed formation of the unstable, unsymmetrical intermediate dimer, {C12O2} or {C1~-C1O2} as
shown in reaction [2] (Emmenegger and Gordon, 1967).
[2]
c/2 + cioi = {Ci-cio2} + cr
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4. CHLORINE DIOXIDE
In some generators that operate with relatively low initial reactant concentrations, a significant
amount of chlorate is formed by reactions with {C12O2}, as shown in reactions [3a], [3b], and [3c].
{ C/202 } + H2O = ClOi + CV + 2H+ [3a]
{ C12O2} + HOCl = CIO{ + Cr + H+ [3b]
{ C12O2 } + 3HOCI + H2O = 2CIO3- + 5H+ + 3Cf [3c]
Highly acidic (pH <3) reaction mixtures force the degradation of {C\2O2} to chlorate rather than
chlorine dioxide, as well as the direct oxidation of chlorite to chlorate.
The overall reactions that describe chlorate ion formation are:
CIO{ + HOCl = CIO3 + Cr + H+ [4a]
and
CIO{ + C/2 + H2O = ClOi + 2Cr + H* [4b]
The following conditions may also produce the chlorate ion:
• Excessively high ratios of C\2 gas:C!O2~.
• Presence of high concentrations of free chlorine at low pH in aqueous solutions.
• Dilute chlorite solutions held at low pH.
• Base-catalyzed disproportionation of chlorine dioxide at high pH values (pH >1 1).
• Reaction mixtures that are highly acidic (pH <3).
• An excess of hypochlorous acid will directly oxidize chlorite ions to chlorate ions rather than
to chlorine dioxide (independent of the rapid formation of the {C^Oa} intermediate).
4.2.4 Generator Design
As hypochlorous acid is formed under acidic conditions, the lowering of optimal concentrations of
precursor reactants will also increase chlorate levels in the generator by promoting reaction [3b].
Therefore, if weak precursor feed stocks or high amounts of dilution water are added to the
generator, chlorate will be more prevalent (according to reaction [3a]). These limitations explain
why generators most often use -25 percent chlorite solutions and gaseous (or near-saturated aqueous)
chlorine. Higher strength solutions of sodium chlorite (e.g., 37 percent) also are more susceptible to
crystallization or stratification at ambient temperatures as high as 25°C(78°F).
Due to these dilution effects, some systems function best as "intermittent batch" generators, (that
produce high concentrations of chlorine dioxide) rather than as "continuous" generators (that produce
lower concentrations (< Ig/L) of chlorine dioxide). The stored solutions are pumped or injected from
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4. CHLORINE DIOXIDE
the storage tank. Cycling frequently avoids long-term (over 24 hour periods) storage of the
generated solution.
Chlorine loop-type systems can obtain high conversion rates if excess chlorine is always present.
Excess chlorine permits the molecular chlorine reaction mechanism (described above) to proceed.
The low pH of the mixture also minimizes the contribution of OH" formed via equation [Ib] by
neutralizing it. These solutions may still be contaminated with excess chlorine needed to drive the
conversion of chlorite ion, but not to the same degree as found in simple aqueous chlorine systems
when operated under dilute conditions. Chlorine-loop generators run best at high capacity since the
chlorite ion is most available in this production mode.
Conventional or acid-enhanced generators produce chlorine dioxide through the intermediate
{CbOi} as long as relatively high concentrations of reactants (-above 20-30 g/L) are maintained in
the reaction chamber prior to dilution. Vapor-phase, recycled loop, and solid chlorite-type generators
that minimize dilute aqueous reaction conditions can obtain high efficiencies by preventing any
chlorite ion from reacting in the "slower" steps described above. This is accomplished by
establishing conditions that force the immediate reaction between chlorite ions and gas-phase or
molecular chlorine at a rate hundreds of times faster than the C\2 hydrolysis in water. This
essentially minimizes the impact of competitive chlorine hydrolysis or acidification on the dominant
Cla gas] mechanism, and prevents the chlorite ion reacting with hypochlorous acid directly.
In all generators, large excess amounts of Cla may result in the over-oxidization of chlorite and
directly form chlorate in aqueous solution (reaction [4b]). Precursor chemical feed rates for the
generators should always be adjusted to chart settings supplied with generators, notably, with the
continuous flow, direct gas injection systems. Re-calibration of these systems is sometimes needed
on-site if feed stock sodium chlorite is not of the correct strength, or if pre-calibrated flow devices
have been replaced.
If aqueous chlorine solutions are mixed with sodium chlorite feed stock solutions, the following
mechanisms are dominant, which may affect the formation rates of chlorine dioxide:
• Chlorine gas reacts with water to form hypochlorous and hydrochloric acids, rather than
directly with chlorite to form chlorine dioxide. (Water and chlorite both compete for the Cla
molecule simultaneously) (see equations [4.1a-c] and Section 6.1.1).
• Chlorate ion is formed (reactions [3a], [3b], and [3c]).
• Only 4 moles of chlorine dioxide are obtained from 5 moles of sodium chlorite via direct
acidification (reaction [lc]). This may become important at low pH and high chloride ion
levels.
The practical side of all of this is that different generators operate under different optimal conditions.
For example, reactor columns should not be continuously flooded with excess water in vapor-phase
systems. It is the main reason why dry chlorite-based generator reactor columns should not get wet.
Over-dilution of the precursor reactants themselves will lower conversion efficiencies due to the
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4. CHLORINE DIOXIDE
favored formation of chlorate over chlorine dioxide. Batch-type generation should always be carried
out at maximal CIC^ concentration with appropriate adjustments at the pump (located downstream of
the reactor at the batch tank) for dosage or flow. Changes in chlorine dioxide concentrations in the
batch tanks would then be minimized, and pump calibration does not need to include a broad range
of chlorine dioxide levels. For the newer gas chlorine generators using dry sodium chlorite in an
inert matrix, small amounts of humidifying water in the mixture do not interfere significantly with
the simple gaseous C12:C1O2~ reaction. These small traces of water allow for continuous exposure of
on the inert surfaces within the reactor column.
Chlorine dioxide generators are relatively simple mixing chambers. The reactors are frequently filled
with media (Teflon™ chips, ceramic or raschig rings) to generate hydraulic turbulence for mixing. A
sample petcock valve on the discharge side of the generator is desirable to allow for monitoring of
the generation process.
The Recommended Standards for Water Works (GLUMRB, 1992) and drinking water design
textbooks such as Unit Processes in Drinking Water Treatment by Masschelein (1992) are excellent
sources for chlorine dioxide generation design criteria and application.
4.2.5 Chemical Feed Systems
Fiberglass Reinforced vinyl ester Plastic (FRP) or High Density Linear Polyethylene (HDLPE) tanks
with no internal insulation or heat probes are recommended for bulk storage of 25 to 38 percent
solution sodium chlorite. Nozzles should include truck unloading vents and local level and
temperature indication. Transfer pumps should be centrifugal with 316 stainless steel, fiberglass,
Hypalon™, wetted Teflon™ parts, or epoxy resins. The pump should be sealless or equipped with
double mechanical seals. The recommended piping material is CPVC, although vinyl ester or Teflon
™ piping systems are acceptable. Carbon steel and stainless steel piping systems are not
recommended.
Depending upon system size, sodium chlorite can be purchased in 55-gallon drums, 275-gallon non-
returnable totes, or in bulk quantities. A 30-day storage supply of sodium chlorite can easily be met
for most small systems by using 55-gallon drums. A 55-gallon drum weighs approximately 600 Ibs.
Equipment should be provided such that one person can easily handle a drum. All gaseous chlorine
or hypochlorite solution-related plumbing should follow Chlorine Institute directives.
Storage and chlorine dioxide systems typically include the following:
• Storage and feeding in a designated space.
• Use of non-combustible materials such as concrete for construction.
• Storage in clean, closed, non-translucent containers. Exposure to sunlight, UV light, or
excessive heat will reduce product strength.
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4. CHLORINE DIOXIDE
* Avoid storage and handling of combustible or reactive materials, such as acids or organic
materials, in the sodium chlorite area.
• Secondary containment for storage and handling areas to handle the worse case spill with
sumps provided to facilitate recovery.
• A water supply near storage and handling areas for cleanup.
• Inert material should be used in contact with the strong oxidizing and/or acid solutions
involved in chlorine dioxide systems.
• Storage tanks with vents to outside.
• Adequate ventilation and air monitoring.
• Gas masks and first aid kits outside of the chemical areas.
• Reactor with glass view ports if it is not made of transparent material.
• Flow monitoring on all chemical feed lines, dilution water lines, and chlorine dioxide
solution lines.
• Dilution water should not be excessively hard in order to avoid calcium deposits and should
be near neutral pH.
• On-site and frequent testing ofxhemical solution strengths should be practiced to achieve
efficient process control.
* Air contact with chlorine dioxide solutions should be controlled to limit the potential for
explosive concentrations possibly building up within the generator. Chlorine dioxide
concentrations in air higher than 8 to 10 percent volume should be avoided. Two methods
can be applied: operation under vacuum or storage under higher positive pressure (45 to 75
psig) to prevent buildup of gas-phase C1O2 in the head space. Bulk storage (batch) tanks
containing C1O2 should be suitably vented to atmosphere.
Sodium chlorite solution feed pumps are commonly diaphragm-metering pumps for liquid feed rate
control. If centrifugal pumps are used, the only acceptable packing material is Teflon. If lubrication
is needed, minimum quantities of fire-resistant lubricants should be used. Pump motors should be
totally enclosed, fan-cooled (TEFC) with sealed-for-life bearings. Couplings should be of the
greaseless type. Water lines for mechanical seals should have a pressure gauge and throttling valve
on the outlet side. Visual means should be provided to verify adequate water flow. Each pump
should include a calibration chamber.
Pipes carrying sodium chlorite should be provided sufficient support to minimize risk of
overstressing joints. Flexible connections to pumps should also be provided to minimize risk of
vibration damage. Pipe should be sloped to drainage points and valved hose connections provided at
strategic points for efficient flushing and draining. Service water for flushing feed lines should be
introduced only through temporarily connected hoses protected by a backflow preventer. Service
water lines should include check valves. Hose connections from service water lines should have a
vent valve to release pressure before the hose is disconnected after use.
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4. CHLORINE DIOXIDE
Flows are frequently measured with magnetic flow meters, mass flow meters, or rotameters for
precise control. Provisions should always be made for back-flow prevention. Sodium chlorite is
extremely reactive, especially in the dry form, and care should be taken to protect against potentially
explosive conditions.
Chlorine dioxide solution concentrations below about 10 g/L will not produce sufficiently high vapor
pressures to present an explosion hazard under most ambient conditions of temperature and pressure.
In water treatment, chlorine dioxide solution concentrations rarely exceed 4 g/L for temperatures less
than 40°C, and treatment levels generally range from 0.1 to 5.0 mg/L. If temperatures exceed 50°C,
storage tanks should be suitably vented due to the higher levels of ClOa possible. As cold
service/potable water is typically used as generator dilution water, these conditions are rarely
encountered.
4.2.6 Generator Power Requirements
Generator power requirements are similar to those for chlorination systems. For all generators (20 to
12,000 Ib/day) power can be supplied from 120 VAC single phase, to 480 VAC three phase. Power
demand will vary based upon make-up water pressure available to operate the venturi. Fractional
horsepower metering pumps are required, based upon system configuration.
4.3 Primary Uses and Points of Application for Chlorine
Dioxide
The calculation of CT for chlorine dioxide is similar to other disinfectants, with accurate
determinations of residual concentrations being a prerequisite for effective disinfection. Primary
disinfectant credit is achieved by the residual concentration and the effective contact time. It has
been found in practice that because of the volatile nature of the gas, chlorine dioxide works
extremely well in plug flow reactors such as pipe lines. It can be easily removed from dilute aqueous
solution by turbulent aeration in rapid mix tanks or purging in recarbonation basins. CT credits
should not be expected through a filter because the likelihood that no residual remains in the filtered
water (DeMers and Renner, 1992). For post CT disinfection credit, chlorine dioxide can be added
before clearwells or transfer pipelines. Ample sampling points should be included to allow close
monitoring of residual concentrations. It is well known that chlorine dioxide is commonly destroyed
by UV in basins exposed to sunlight or bright fluorescent lights. Therefore, protective design
elements should be incorporated if such exposure is anticipated.
4.3.1 Disinfection
Before chlorine dioxide is selected for use as a primary disinfectant an oxidant demand study should
be completed. Ideally, this study should consider the seasonal variations in water quality,
temperature, and application points. Table 4-2 shows typical results for a single sample of a demand
study completed on a surface water source.
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The MRDL for chlorine dioxide is 0.8 mg/L and the MCL for chlorite is 1.0 mg/L per the D/DBP
rule. This means that if the oxidant demand is greater than about 1.4 mg/L, chlorine dioxide may not
be used as a disinfectant because the chlorite/chlorate ions byproduct, might exceed the maximum
level allowed, unless inorganic byproducts (e.g., chlorite) are subsequently removed. There are
numerous means to reduce excessive chlorite levels prior to chlorination in conventional water
plants.
Table 4-2. Surface Water Chlorine Dioxide Demand Study Results
Dose (mg/L)
1.4
Time (min)
3
10
20
40
60
CI02(mg/L)
0.47
0.30
0.23
0.16
0.11
CI02-(mg/L)
0.76
0.98
1.08
1.11
1.11
CI03-(mg/L)
0.05
0.06
0.07
0.07
0.07
Source: DeMers and Renner, 1992.
Note: 'Raw water sample, 23°C, 8.5 pH.
Typical dosages of chlorine dioxide used as a disinfectant in drinking water treatment range from
0.07 to 2.0 mg/L. For plants using chlorine dioxide, median concentrations of chlorite and chlorate
were found to be 0.24 and 0.20 mg/L, respectively in an EPA survey (USEPA, 1998), the standard is
1.0 mg/L.
4.3.2 Taste and Odor Control
A common application of chlorine dioxide in drinking water in the United States has been for control
of tastes and odors associated with algae and decaying vegetation. Chlorine dioxide is also effective
in destroying taste and odor producing phenolic compounds. The recommended location for
application of chlorine dioxide for this purpose will depend on raw water quality, the type of
treatment plant and any other purposes for chlorine dioxide addition. In conventional treatment
plants, it is recommended that chlorine dioxide be added near the end of or following, the
sedimentation basin. If the raw water turbidity is low (for example, less than 10 NTU), chlorine
dioxide can be added at the beginning of the plant. Some utilities follow this practice because
chlorine dioxide is effective in controlling algae growth in flocculation and sedimentation basins that
are exposed to sunlight (DeMers and Renner, 1992). Such application during periods of darkness
may be more successful for nuisance algae control.
4.3.3 Oxidation of Iron and Manganese
Chlorine dioxide can be used to oxidize both iron and manganese. Chlorine dioxide reacts with the
soluble forms of iron and manganese to form precipitates that can be removed through sedimentation
and filtration. Chlorine dioxide reduces to chlorite ion in this reaction (Knocke et al., 1993). About
1.2 mg/L of chlorine dioxide is required to remove 1.0 mg/L of iron, and 2.5 mg/L of chlorine
dioxide are required to remove 1.0 mg/L of manganese. For high concentrations of iron and
manganese, the use of chlorine dioxide is limited to the 1.0 mg/L chlorite/chlorate ion byproduct, as
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4. CHLORINE DIOXIDE
described before. Ferrous iron may be added prior to conventional coagulation to chemically reduce
chlorite ion (to chloride ion) and improve the overall flocculation process.
4.4 Pathogen Inactivation and Disinfection Efficacy
For water treatment, chlorine dioxide has several advantages over chlorine and other disinfectants.
In contrast to chlorine, chlorine dioxide remains in its molecular form in the pH range typically found
in natural waters (Roberts et al., 1980). Chlorine dioxide is a strong oxidant and disinfectant. Its
disinfection mechanisms are not well understood, but appear to vary by type of microorganism
4.4.1 Inactivation Mechanisms
Gross physical damage to bacterial cells or viral capsids has not been observed at the low
concentrations of chlorine dioxide typically used to disinfect drinking water. Therefore, studies have
focused primarily on two more subtle mechanisms that lead to the inactivation of microorganisms:
determining specific chemical reactions between chlorine dioxide and biomolecules; and observing
the effect chlorine dioxide has on physiological functions.
In the first disinfection mechanism, chlorine dioxide reacts readily with amino acids cysteine,
tryptophan, and tyrosine, but not with viral ribonucleic acid (RNA) (Noss et al., 1983; Olivier et al.,
1985). From this research, it was concluded that chlorine dioxide inactivated viruses by altering the
viral capsid proteins. However, chlorine dioxide reacts with poliovirus RNA and impairs RNA
synthesis (Alvarez and O'Brien, 1982). It has also been shown that chlorine dioxide reacts with free
fatty acids (Ghandbari et al., 1983). At this time, it is unclear whether the primary mode of
inactivation for chlorine dioxide lies in the peripheral structures or nucleic acids. Perhaps reactions in
both regions contribute to pathogen inactivation.
The second type of disinfection mechanism focuses on the effect of chlorine dioxide on physiological
functions. It has been suggested that the primary mechanism for inactivation was the disruption of
protein synthesis (Bernarde et al., 1967a). However, later studies reported the inhibition of protein
synthesis may not be the primary inactivation mechanism (Roller et al., 1980). A more recent study
reported that chlorine dioxide disrupted the permeability of the outer membrane (Aieta and Berg,
1986). The results of this study were supported by the findings of Olivieri et al. (1985) and
Ghandbari et al. (1983), which found that the outer membrane proteins and lipids were sufficiently
altered by chlorine dioxide to increase permeability.
4.4.2 Environmental Effects
Studies have been performed to determine the effect of pH, temperature, and suspended matter on the
disinfection efficiency of chlorine dioxide. Following is a summary of the effects these parameters
have on pathogen inactivation.
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4.4.2.1 pH
In comparison to chlorine, studies have shown that pH has much less effect on pathogen inactivation
for viruses and cysts with chlorine dioxide than with chlorine in the pH range of 6 to 8.5. Unlike
chlorine, studies on chlorine dioxide have shown the degree of inactivation of poliovirus 1 (Scarpino
et al., 1979) and Naegleria gruberi cysts (Chen et al., 1984) increase as the pH increases.
The results of studies on E. coli inactivation are inconclusive. It has been found that the degree of
inactivation by chlorine dioxide increases as pH increases (Bernarde et al., 1967a). However, an
earlier study found that the bactericidal activity of chlorine dioxide was not affected by pH values in
the range of 6.0 to 10.0 (Ridenour and Ingols, 1947). A recent study on Cryptosporidium found that
inactivation of oocysts using chlorine dioxide occurred more rapidly at a pH of 8.0 than 6.0. At a
similar CT value, the level of inactivation at pH of 8.0 was approximately twice that at a pH of 6.0
(Le Chevallier et al., 1997). Another study found that chlorine dioxide efficacy increases for Giardia
inactivation at higher pH levels and that this may be the result of chemical or physical changes in
Giardia cyst structure rather than pH effects on chlorine dioxide disproportionation (Liyanage et al.,
1997). More research is needed to further clarify how pH impacts the effectiveness of chlorine
dioxide.
4.4.2.2 Temperature
Similar to chlorine, the disinfection efficiency of chlorine dioxide decreases as temperature decreases
(Ridenour and Ingols, 1947). This finding is supported by the data from Chen et al. (1984) shown in
Figure 4-3 for the inactivation of Naegleria gruberi cysts. The curve shows the CT required to
achieve 99 percent inactivation for temperatures between 5 and 30°C.
In a more recent study, LeChevallier et al. (1997) found that reducing the temperature from 20°C to
10°C reduced the disinfection effectiveness of chlorine dioxide on Cryptosporidium by 40 percent,
which is similar to previous results for Giardia and viruses. Gregory et al. (1998) found that even
under the most favorable conditions (i.e., at a pH of 8.5), required doses to achieve 2-log
Cryptosporidium inactivation do not appear to be a feasible alternative, requiring doses of more than
3.0 mg/L with a 60 minute detention time. At neutral pH levels, the required doses may be more
than 20 mg/L.
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4. CHLORINE DIOXIDE
15 20
Temperature (°C)
25
Figure 4-3. Effect of Temperature on N. GruberiCyst Inactivation at pH 7
4.4.2.3 Suspended Matter
Suspended matter and pathogen aggregation affect the disinfection efficiency of chlorine dioxide.
Protection from chlorine dioxide inactivation due to bentonite was determined to be approximately
11 percent for turbidities equal to or less than 5 NTUs and 25 percent for turbidities between 5 and
17NTUs(Chenetal., 1984).
Laboratory studies of poliovirus 1 preparations containing mostly viral aggregates took 2.7 times
longer to inactivate with chlorine dioxide than single state viruses (Brigano et al., 1978). Chen et al.
(1984) also found that clumps ofNaegleria gruberi cysts -were more resistant to chlorine dioxide than
unclumped cysts or clumps of smaller size.
4.4.3 Disinfection Efficacy
Several investigations have been made to determine the germicidal efficiency of chlorine dioxide
since its introduction in 1944, as a drinking water disinfectant. Most of the investigations were
carried out as a comparison to chlorine; some studies have compared chlorine dioxide and ozone.
Chloride dioxide is a more effective disinfectant than chlorine but is less effective than ozone.
4.4.3.1 Bacteria Inactivation
Quantitative data were published as early as the 1940s demonstrating the efficacy of chlorine dioxide
as a bactericide. In general, chlorine dioxide has been determined to be equal to or superior to
chlorine on a mass-dose basis. It was demonstrated that even in the presence of suspended matter,
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chlorine dioxide was effective against E. coli and Bacillus anthracoides at dosages in the range of 1
to 5 mg/L (Trakhtman, 1949). Ridenour and Armbruster (1949) reported that an orthotolidine
arsenite (OTA) chlorine dioxide residual of less than 1 mg/L was effective against Eberthella
typhosa, Shigella dysenteriae, and Salmonella paratyphi B. Under similar pH and temperature
slightly greater OTA residuals were required for the inactivation of Pseudomonas aeruginosa and
Staphylococcus aitreus.
Chlorine dioxide was shown to be more effective than chlorine at inactivating B. subtilis, B.
mesentericus, and B. megatherium spores (Ridenour et al., 1949). Moreover, chlorine dioxide was
shown to be just as effective or more effective than chlorine at inactivating Salmonella typhosa and
S. paratyphi (Bedulivich et al., 1954).
In the early 1960s several important contributions were made by Bernarde et al. (1967a and 1967b).
Chlorine dioxide was found to be more effective than chlorine at disinfecting sewage effluent and the
rate of inactivation was found to be rapid.
A comprehensive investigation of chlorine dioxide as disinfectant was performed by Roberts et al.
(1980). The investigation was performed using secondary effluents from three different wastewater
treatment plants. One of the objectives was to determine the relationships between dosages and
contact times and bactericidal efficiency. Dosages were compared for 2, 5, and 10 mg/L of chlorine
dioxide and chlorine. The contact times selected were 5, 15 and 30 minutes. Results of the
investigation are shown in Figure 4-4. As shown, chlorine dioxide demonstrated a more rapid
coliform inactivation than chlorine at the shortest contact time of 5 minutes and higher
concentrations. However, after 30 minutes of contact time, chlorine dioxide was equal or slightly
less efficient than chlorine as a bactericide.
Oliveri et al. (1984) studied the effectiveness of chlorine dioxide (and chlorine) residuals in
inactivating total coliform and f2 coliphage virus in sewage introduced to a water distribution system.
Initial chlorine dioxide residuals between 0.85 and 0.95 mg/L resulted in an average 2.8-log
inactivation of the total coliform and an average 4.4-log inactivation of the f2 coliphage virus, over a
contact time of 240 minutes.
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4. CHLORINE D/OX/DE
3
o
Ct
o
o:
o
LL
O
O
§
2 mg/i
5 mg/t
-5
-6
-71—
LEGEND
0 D C102
• • CI2
}1O mg/t
15
CONTACT TIME (min)
30
Source: Roberts et al., 1980.
Figure 4-4. Comparison of Germicidal Efficiency of Chlorine Dioxide and Chlorine
4.4.3.2 Protozoa Inactivation
The disinfection efficiency of chlorine dioxide has been shown to be equal to or greater than chlorine
for Giardia inactivation. Based on a 60 minute contact time, chlorine dioxide doses in the range of
1.5 to 2 mg/L are capable of providing a 3-log Giardia inactivation at 1°C to 25°C and pHs of 6 and
9 (Hofmann et al., 1997). Depending on the temperature and pH, Cryptosporidium has been found to
be 8 to 16 times more resistant to chlorine dioxide than Giardia (Hofmann et al., 1997). Although
some Cryptosporidium oocysts remained viable, one group of researchers found that a 30-minute
contact time with 0.22 mg/L chlorine dioxide could significantly reduce oocyst infectivity (Peeters et
al., 1989). In contrast, other researchers have found that CT values in the range of 60 to
80 mg-min/L were necessary to provide 1- to 1.5-log inactivation (Korich et al., 1990; Ransome et
al., 1993). Finch et al. (1995) reported that the CT values for 1-log inactivation was in the range of
27 to 30 mg-min/L. For 2-log inactivation, the CT value was approximately 40 mg-min/L, and 70
mg-min/L for 3-log inactivation. Finch et al. (1997) found 3-log inactivation of Cryptosporidium
oocysts with initial chlorine dioxide residual concentrations of 2.7 and 3.3 mg/L for contact times of
120 minutes, at pH of 8.0 and a temperature of 22°C.
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Both Chen et al. (1985) and Sproul et al. (1983) have investigated the inactivation ofNaegleria
gruberi cysts by chlorine dioxide. Both studies concluded that chlorine dioxide is an excellent
disinfectant against cysts and that chlorine dioxide is better than or equal to chlorine in terms of
inactivation. Chlorine dioxide was found to be superior to chlorine at higher pHs. However, the
authors cautioned that the CT required for 2-log inactivation was much higher than normally
employed for water treatment at that time.
4.4.3.3 Virus Inactivation
Chlorine dioxide has been shown to be an effective viricide. Laboratory studies have shown that
inactivation efficiency improves when viruses are in a single state rather than clumped. It was
reported in 1946 that chlorine dioxide inactivated Poliomyelitis (Ridenour and Ingols, 1946). This
investigation also showed that chlorine dioxide and free chlorine yielded similar results. Other
studies have verified these findings for poliovirus 1 (Cronier et al., 1978) and Coxsackie virus A9
(Scarpino, 1979). At greater than neutral pHs (where hypochlorite ion is the predominant species)
chlorine dioxide has been found to be superior to chlorine in the inactivation of numerous viruses
such as echovirus 7, coxsackie virus B3, and sendaivirus (Smith and McVey, 1973). Sobsey (1998)
determined CT values based on a study of Hepatitis A virus, strain HM 175. The study found 4-1 og
inactivation levels are obtainable at CT values of less than 35 at 5°C and less than 10 at a temperature
of25°C.
4.4.3.4 CT Values
Chlorine dioxide is regarded as a strong disinfectant that is effective at inactivating bacterial, viral,
and protozoan pathogens. CT values for Giardia and virus inactivation are shown in Figure 4-5 and
Figure 4-6, respectively (AWWA, 1991).
CT values shown in Figure 4-5 are based on disinfection studies using in vitro excystation of Giardia
muris. Average CT values for 2 log removal were extrapolated using first order kinetics and
multiplied by a safety factor of 1.5 to obtain the CT values for other log removal CT values. Due to
the limited amount of data available at pH values other than 7, the same CT values are used for all
pHs. Because chlorine dioxide is more effective at a pH 9 than at a pH of 7, the CT values shown in
Figure 4-5 are more conservative for higher pHs than for lower pHs. A lower safety factor was used
to derive the CT values for chlorine dioxide than for ozone due to the fact that the chlorine dioxide
values were derived from Giardia muris studies, which are more resistant than Giardia lamblia.
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4. CHLORINE DIOXIDE
25.0 '
20.0 ''
-0.5-log Inactivation
~ 1 -log Inactivation
~ 1.5-!og Inactivation
~2-log Inactivation
-2.5-log Inaotivation
-3-log Inactivation
10.0 ••
15
Temperature (°C)
Source: AWWA, 1991.
Figure 4-5. CT Values for Inactivation of Giardia Cysts by Chlorine Dioxide
CT values shown in Figure 4-6 were obtained by applying a safety factor 2 to the average CT values
derived from the studies on hepatitis A virus, strain HM 175 (Sobsey, 1988). CT values at
temperatures other than 5°C were derived by applying a twofold decrease for every 10°C increase in
temperature.
Figure 4-7 and Figure 4-8 show the relationship between CT products and log inactivation of
Cryptosporidium at 20 and 10°C, respectively, and pHs of 6 and 8. CT values shown in Figure 4-7
and Figure 4-8 indicate that oocysts were more rapidly inactivated at pH 8 than 6 and that
temperature does impact the disinfection efficiency of chlorine dioxide. Reducing the temperature
from 20 to 10°C reduced the disinfection effectiveness by 40 percent. Finch (1997) is studying
Cryptosporidium inactivation under laboratory conditions using a variety of different disinfectants,
one of which is chloride dioxide.
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35,0
2-log Inactivation
3-log Inactivation
4-log Inactivation
15
Temperature (°C)
Source: AWWA, 1991,
Figure 4-6. CT Values for Inactivation of Viruses by Chlorine Dioxide
4.5 Chlorine Dioxide Disinfection Byproducts
Byproducts from the use of chlorine dioxide include chlorite, chlorate, and organic DBFs. This
section discusses the formation of these byproducts and methods to reduce or remove these DBFs.
The use of chlorine dioxide aids in reducing the formation of TTHMs and HAAs by oxidizing
precursors, and by allowing the point of chlorination to be moved farther downstream in the plant
after coagulation, sedimentation, and filtration have reduced the quantity of NOM.
4.5.1 Production of Chlorite and Chlorate
Chlorite and chlorate are produced in varying ratios as endproducts during chlorine dioxide treatment
and subsequent degradation. The primary factors affecting the concentrations of chlorine dioxide,
chlorite, and chlorate in finished drinking water involve:
• Dosage applied/oxidant demand ratio.
• Blending ratios of sodium chlorite and chlorine during the chlorine dioxide generation
process.
• Exposure of water containing chlorine dioxide to sunlight.
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4. CHLORINE DIOXIDE
1.25"
.S
I
0.75 '
0.25 ''
40 50
CT Product (mg min/L)
80
90
- pH 6.0; 1.52 mg/L dose, 1.23 mg/L residual
-pH 6.0; 0.51 mg/L dose, 0.38 mg/L residual
- pH 8.0:1.52 mg/L dose, 1.23 mg/L residual
~pH 8.0; 0.51 mg/L dose, 0.39 mg/L residual
Source: LeChevallier et al., 1996.
Figure 4-7. C. parvum Inactivation by Chlorine Dioxide at 20°C
0.75 ''
3
0.25 -'
20
30
40 50
CT Product (mg min/L)
60
70
80
90
- pH 6.0; 1.52 mg/L dose, 1.23 mg/L residual
-pH 6.0; 0.51 mg/L dose. 0.39 mg/L residual
- pH 8.0; 1.52 mg/L dose, 1.23 mg/L residual
~pH 8.0; 0.51 mg/L dose, 0.39 mg/L residual
Source: LeChevallier et al., 1996.
Figure 4-8. C. parvum Inactivation by Chlorine Dioxide at 10°C
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• Reactions between chlorine and chlorite if free chlorine is used for distribution system
residual maintenance.
• Levels of chlorate in sodium chlorite feedstock.
Incomplete reaction or non-stoichiometric addition of the sodium chlorite and chlorine reactants can
result in unreacted chlorite in the chlorine dioxide feed stream. Dilute chlorine dioxide solutions are
stable under low or zero oxidant-demand conditions. The quantity of chlorate produced during the
chlorine dioxide generation process is greater with excess chlorine addition. Likewise, a low or high
pH can increase the quantity of chlorate during the chlorine dioxide generation process. See Section
4.2, "Generation," for a detailed discussion of the chemistry of chlorine dioxide generation.
Numerous inorganic and biological materials found in raw water will react with chlorine dioxide
(Noack and Doerr, 1977). Chloride (Cl") and chlorite (C1O2~) ions are the dominant degradation
species arising from these reactions, although chlorate (C1O3~) can appear for a variety of reasons
when chlorine dioxide is used (Gordon et al., 1990; Werdehoff and Singer, 1987). The immediate
redox reactions with natural organic matter play the dominant role in decay of chlorine dioxide into
chlorite in drinking water (Werdehoff and Singer, 1987). Chlorite ion is generally the primary
product of chlorine dioxide reduction. The distribution of chlorite and chlorate is influenced by pH
and sunlight. Approximately 50 to 70 percent of the chlorine dioxide consumed by oxidation
reactions is converted to chlorite under conditions typical in water treatment (Rav-Acha et al., 1984;
Werdehoff and Singer, 1987). The application of 2 mg/L chlorine dioxide is expected to produce 1
to 1.4 mg/L of chlorite (Singer, 1992).
Chlorite is relatively stable in the presence of organic material but can be oxidized to chlorate by free
chlorine if added as a secondary disinfectant (Singer and O'Neil, 1987).
aoi + ocr = cioi + cr
Chlorate is therefore produced through the reaction of residual chlorite and free chlorine during
secondary disinfection.
In addition, chlorine dioxide also disproportionates under highly alkaline conditions (pH>9) to
chlorite and chlorate according to the following reaction:
2C/02 + 2OH- = CIO2~ + CIO3- + H2O
In water treatment processes that require high pH, such as softening, chlorine dioxide should be
added after the pH has been lowered (Aieta et al., 1984).
The occurrence of photochemical decomposition of chlorine dioxide can affect the ultimate
concentrations of chlorine dioxide, chlorite, and chlorate in water treated with chlorine dioxide.
Moreover, generally, sunlight may increase chlorate concentrations in uncovered storage basins
containing water with chlorine dioxide residuals. Exposure to ultraviolet light will also change the
potential reactions between chlorine dioxide and the bromide ion.
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4.5.2 Organic DBFs Produced by Chlorine Dioxide
Chlorine dioxide generally produces few organic DBFs. However, Singer (1992) noted that the
formation of non-halogenated organic byproducts of chlorine dioxide has not been adequately
researched, and expected that chlorine dioxide will produce the same types of oxidation byproducts
that are produced through ozonation. The application of chlorine dioxide does not produce THMs
and produces only a small amount of total organic halide (TOX) (Werdehoff and Singer, 1987).
A study was conducted in 1994 by Richardson et al., to identify semivolatile, organic DBFs produced
by chlorine dioxide treatment in drinking water. Samples were taken from a pilot plant in Evansville,
Indiana that included the following treatment variations:
• Aqueous chlorine dioxide;
• Aqueous chlorine dioxide, ferrous chloride, (FeC^), chlorine (Ck), and dual media filtration
(sand and anthracite);
• Gaseous chlorine dioxide; and
• Gaseous chlorine dioxide, ferrous chloride (FeCli), chlorine (Ck), and dual media filtration
(sand and anthracite).
Using multispectral identification techniques, more than 40 different DBFs (many at sub-nanogram/L
[ng/L] levels) were identified including carboxylic acids and maleic anhydrides isolated from
XAD™ concentrates, some of which may be regulated in the Stage 2 DBPR. THMs were not found
after chlorine dioxide was added to the water; however, THMs did show up during subsequent
chlorination.
4.5.3 Chlorine Dioxide DBP Control Strategies
EPA recommends that the total concentration of chlorine dioxide, chlorite, and chlorate be less than
1.0 mg/L as C\2 (USEPA, 1983). In addition, chlorine dioxide concentrations exceeding 0.4 to 0.5
mg/L contribute to taste and odor problems (AWWA, 1990). Due to these issues, the use of chlorine
dioxide to provide a disinfectant residual is somewhat limited in moderate to high TOC water. In
low oxidant-demand water, however, C1O2 residuals may last several days.
Once formed, chlorate is stable in finished drinking water. No known treatment exists for removing
chlorate once it is formed. However, three strategies (Gallagher et al., 1994) that have been proven
effective for chlorite removal are:
• Adding reduced-sulfur compounds such as sulfur dioxide and sodium sulfite (not
recommended).
• Applying either granular activated carbon (GAC) or powdered activated carbon (PAC).
• Adding reduced iron salts, such as ferrous chloride and ferrous sulfate.
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Chlorite removal from drinking water through sulfur dioxide and other sulfur-based reducing agents
has been reported effective, but not desirable. A study of chlorite removal by sulfur dioxide indicates
that a lower pH level yields higher chlorite removal, and chlorite removal efficiencies increase as the
sulfur dioxide dose increases. Unfortunately, this removal process forms significant levels of
chlorate when sulfur dioxide and metasulfite are utilized. Therefore, it is concluded that treatment
with sulfur dioxide and metasulfite is not desirable for chlorite removal (Dixon and Lee, 1991). In
addition, sodium thiosulfate results in effective chlorite reduction, but the degree of removal is highly
dependent upon pH and contact time and relatively high dosages are required. Again, this
application of sodium thiosulfate is not desirable because the required dosages are too high (Griese et
al., 1991).
The addition of ferrous iron in drinking water is effective for chlorite removal, with chloride the
expected byproduct. Chlorite reduction occurs quickly in the pH range of 5 to 7, and complete
reduction occurs within 3 to 5 seconds. Excess reduced iron remaining in solution reacts with
dissolved oxygen at neutral pH, but under acidic conditions (pH < 6.5) the stability of the soluble
iron can create aesthetic (staining) problems if excess iron is used. Special consideration should be
given to ferrous iron dosage requirements so that the secondary MCL for iron is not exceeded
(Knocke and latrou, 1993).
Chlorite can be controlled by PAC at relatively high dosages (10 to 20 mg/L) and low pHs (5.5 to
6.5). Unless PAC is used for other purposes, such as odor control, it requires large doses and is not
cost effective. PAC brands can differ in their capacity to reduce chlorite.
GAC can remove chlorite but breakthrough may occur relatively early. The removal of chlorite by
GAC appears to be a result of adsorption and chemical reduction (Dixon and Lee, 1991). There is an
initial high removal efficiency due to chlorite adsorption. As the adsorptive sites are occupied,
chemical reduction on the GAC surface becomes the primary removal mechanism. This results in an
initial high removal efficiency. Although chlorite levels exiting the GAC filters are low, the chlorate
levels are high, most likely a result of reactions in the GAC filters between chlorite and free chlorine.
According to studies, the capacity of GAC beds is low, and if free chlorine and chlorite ion are
present in the GAC influent, chlorate ion will form. The most effective way to operate GAC for
chlorite reduction and avoid chlorate is to minimize production run times and have no chlorine
present in the filter.
4.6 Status of Analytical Methods
In addition to the monitoring requirements that apply regardless of the disinfectant used, the DBPR
requires that water systems that use chlorine dioxide for disinfection or oxidation must also monitor their
system for chlorine dioxide and chlorite.
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4.6.1 Chlorine Dioxide and Chlorite Analytical Methods
For compliance monitoring for chlorine dioxide, systems must use one of the two methods specified in 40
CFR §141.131(c), including (1) DPD, Standard Method 4500-CLO? D, or (2) Amperometric Method II,
Standard Method 4500-CLO2 E. Where approved by the state, systems may also measure residual
disinfectant concentrations for chlorine dioxide by using DPD colorimetric test kits.
For compliance monitoring for chlorite, systems must use one of the three methods specified in 40 CFR
§141.131(b), including (1) Amperometric Titration, Standard Method 4500-CLO2 E, (2) Ion
Chromatography, EPA Method 300.0, or (3) Ion Chromatography, EPA Method 300.1. The regulations
specify that Amperometric Titration may be used for routine daily monitoring of chlorite at the entrance
to the distribution system, but that Ion Chromatography must be used for routine, monitoring of chlorate
and monthly additional monitoring of chlorate in the distribution system.
Details of these analytical procedures can be found in:
Standard Methods for the Examination of Water and Wastewater, 19th Edition, American Public
Health Association, 1995.
Methods for the Determination of Inorganic Substances in Environmental Samples. USEPA.
1993. EPA/600/R-93/100.
USEPA Method 300.1, Determination of Inorganic Anions in Drinking Water by Ion
Chromatography, Revision 1.0. USEPA. 1997. EPA/600/R-98/118.
Table 4-3 summarizes the analytical methods approved for use for chlorine dioxide and chlorite and
provides some background information for each method.
4.6.2 Chlorine Dioxide Monitoring for Systems Using Chlorine
Dioxide
For chlorine dioxide monitoring, community, non-transient non-community, and transient non-
community water systems that use chlorine dioxide for disinfection or oxidation, are required to take daily
samples at the entrance to the distribution system. For any daily sample that exceeds the chlorine dioxide
MRDL of 0.8 mg/L, the system must take additional samples in the distribution system the following day
at the locations specified in the DBPR, in addition to the daily sample required at the entrance to the
distribution system.
Additional sampling is to be performed in one of two ways, depending on the disinfectant that is used to
maintain a disinfectant residual in the distribution system. If chlorine dioxide or chloramines are used to
maintain a disinfectant residual, or if chlorine is used to maintain the residual and there are no disinfection
addition points after the entrance to the distribution system (i.e., no booster chlorination), the system must
take three samples as close to the first customer as possible, at intervals of at least six hours. If chlorine is
used to maintain a disinfectant residual and there are one or more disinfection addition points after the
entrance to the distribution system, the system must take one sample at each of the following locations:
(1) as close to the first customer as possible, (2) in a location representative of average residence time,
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and (3) as close to the end of the distribution system as possible (reflecting maximum residence time in
the distribution system). Chlorine dioxide monitoring may not be reduced.
If any daily sample taken at the entrance to the distribution system exceeds the MRDL, and on the
following day one (or more) of the three samples taken in the distribution system exceed the MRDL, the
system is in violation of the MRDL. The system must take immediate corrective action to lower the level
of chlorine dioxide below the MRDL, and must notify the public of the acute violation pursuant to 40
CFR § 141.32. The system must also report to the State pursuant to 40 CFR § 141.134.
If any two consecutive daily samples taken at the entrance to the distribution system exceed the MRDL,
the system is also in violation of the MRDL and must notify the public of the non-acute violation
pursuant to 40 CFR §141.32. The system must also report to the State pursuant to 40 CFR §141.134.
4.6.3 Chlorite Monitoring for Systems Using Chlorine Dioxide
For chlorite monitoring, community and non-transient non-community water systems that use chlorine
dioxide for disinfection or oxidation are required to take daily samples at the entrance to the distribution
system. For any daily sample that exceeds the chlorite MCL of 1.0 mg/L, the system must take additional
samples in the distribution system the following day at the locations specified in the DBPR. These
additional samples are to be collected at: (1) a location as close to the first customer as possible, (2) a
location representative of average residence time, and (3) a location as close to the end of the distribution
system as possible (reflecting maximum residence time in the distribution system).
In addition, systems using chlorine dioxide must take a three-sample set each month in the distribution
system similar to the three locations required if the chlorite MCL is exceeded in the sample collected at
the entrance to the distribution system. Specifically, these three-sample sets are to be collected: (1) in a
location near the first customer, (2) in a location representative of average residence time, and (3) at a
location reflecting maximum residence time in the distribution system. Any additional routine sampling
must be conducted in the same three-sample sets at the specified locations. This monthly sampling
requirement may be reduced to quarterly after one year of monitoring where: (1) no individual chlorite
sample taken in the distribution system has exceeded the MCL and (2) the system has not been required to
conduct follow-up monitoring as a result of a daily sample collected at the entrance to the distribution
system. These systems can remain on an annual schedule until either the daily sample or any of the three
individual quarterly samples exceed the MCL, at which time, the system must revert to monthly
monitoring.
If the arithmetic average of any three-sample set exceeds the chlorite MCL of 1.0 mg/L, the system is in
violation of the MCL and must notify the public pursuant to 40 CFR §141.32, in addition to reporting to
the State pursuant to 40 CFR §141.134.
4.7 Operational Considerations
As with all disinfectant selections, the primacy agency should be consulted when selecting
disinfectants. Certain states have their own operational, maintenance, and monitoring requirements
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for the application of chlorine dioxide. California prohibits the use of chlorine dioxide in ground
water systems, according to Merkle et al., 1997. Also, in Texas, the Texas Natural Resources
Conservation Commission (TNRCC) requires the public water supply to sign a bilateral agreement
which outlines a detailed operator qualifications requirement, testing methods, and procedures,
monitoring locations, testing frequency and reporting procedures. The chlorine dioxide
concentration leaving the water treatment plant must be less than 0.8 mg/L and the chlorite
concentration in the distribution system must be less than 1.0 mg/L.
State requirements must be reviewed to determine the cost-effectiveness of utilizing chlorine dioxide
as part of the overall water treatment scheme. Analytical testing and reporting requirements may
have significant labor and cost impacts.
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Table 4-3. Analytical Methods for Chlorine Dioxide and Related Compounds
Method
DPD as Test Kits
Colorimetric
(SM-4500-CI02.G)
DPD-giycine Method
Colorimetrtc
(SM4500-CI02D)
DPD-FAS
Titrimetrte method
(SM4500-CIC-2.D)
5-Step Amperometric
Method 4500-CI02.E
Ion Chromatography
(EPA Method 300.0 or 300.1)
Conductivity
Two-step Amperometric
Method 4500-CI02.E
Basis
Colored oxidation product.
Use of color comparator is not
recommended. Use instrument
detection.
Colored product, free C\z is
masked with glycine as
chloraminacetic acid.
DPD color titration with standard
FAS until red color is
discharged.
I- oxidation; pH control and gas
purging steps. Skilled analyst
needed.
Must use AS9 column, ext.
standards & suppression.
I- Oxidation; pH control.
Amendable to operator-based
dosage control.
Practical method.
Interferences
Mn2+, other Cfc, related
oxidants.
ClCr slowly; other oxidants.
Iron, other oxidants.
Suitable for ClOa generated
solution.
Low levels not okay.
No other oxidants.
Chloramines, CI02; OCI- &
HOCL undetectable.
Cua+, Mn2+, NCr
Accounts for free Cl2, NhfeCI,
CI02' species.
Limits
' > 0.1 mg/L
> 0.1 mg/L
> 0.1 mg/L
- PQL CIC-2-:
0.1 -0.0.5 mg/L;
CI03- at 0.5 mg/L
~ 0.05 mg/L
> 0.1 mg/L,
not ClOs-
Source: Gates, 1998.
Notes: SM = Standard Methods
4.7.1 Process Considerations
The basic components of chlorine dioxide generation systems include:
• Aqueous hypochlorite solution storage and feed system;
« Sodium chlorite storage and feed system;
• Acid storage and feed system (for Direct-Acidification generators);
• Chlorine storage and feed system;
• Chlorine dioxide generator; and
* Chlorine dioxide feed piping and dispersion equipment.
Sodium chlorite storage and feed systems are basically liquid systems that consist of a storage tank(s)
and solution feed pumps. Outside storage of 25 percent solutions (or greater) of sodium chlorite is
not recommended in cold climates since stratification may occur below 4°C (40°F). Any ice
formation may also damage the storage tanks. In some cases, storage might be separated into bulk
tanks and smaller operational or day tanks that are filled periodically. Storage of dark drums for long
periods in hot climates should be avoided since sodium chlorite decomposition will occur. In the
storage area, light fixtures, switches, wiring, and conduit runs should be located to avoid the risk of
sodium chlorite spilling on them.
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4. CHL OPINE DIOXIDE
4.7.2 Generator Operation
A manual chlorine dioxide feed system may be used where the chlorine dioxide dose remains fairly
constant. The reagent chemicals are manually set for the desired chlorine dioxide capacity at a ratio
of chemicals optimized for maximum chlorine dioxide yield. Some generating systems can produce
95 percent pure chlorine dioxide solutions at full design capacity, but purity can vary when the feed
rate is changed. Turndown capacity may be limited by precision of the flow metering devices,
typically 20 percent of rate capacity. Purity can vary when the feed rate is changed significantly.
Feed water alkalinity, operating conditions, and pH also can affect yield. The ratio of reagent
chemicals should be routinely adjusted for optimum operation. Chlorine dioxide generators can be
provided with automated control to provide modulation of chlorine dioxide feed rates based upon
changes in flow (flow paced) and chlorine dioxide demand (residual control). The automatic
modulation of the generators to meet a demand setpoint varies with manufacturer. Generally,
vacuum and combination systems are limited by the hydraulic requirements of the venturi and the
optimum reaction conditions for chlorine dioxide generation. A chemical metering pump or injector
system is then used with a batch production system to control the applied dose of chlorine dioxide.
4.7.3 Feed Chemicals
Chlorine dioxide is generated when sodium chlorite is either oxidized or acidified, or both, under
controlled pH and temperature conditions. Commonly, solutions of 25 percent active sodium chlorite
or less are used in chlorine dioxide generators. The major safety concern for solutions of sodium
chlorite is the unintentional and uncontrollable release of high levels of chlorine dioxide. Such levels
may approach detonation or conflagration concentrations by accidental acidification.
The feedstock acid used by some of the generators is only one source of accidental chemical
acidification. Accidental mixing with large amounts of any reducing agent or oxidizable material
(such as powdered activated carbon or flammable solvents) also represents a significant hazard. The
AWWA Standard B303-95 (a) includes an outline of some of these materials (AWWA, 1995).
Another concern when handling and storing sodium chlorite solutions is crystallization, which occurs
as a result of lower temperatures and/or higher concentrations. Crystallization will plug pipelines,
valves, and other equipment. Sodium chlorite solution should not be allowed to evaporate to a
powder. If dried, this product becomes a fire hazard and can ignite in contact with combustible
materials. A sodium chlorite fire may result in a steam explosion if too much water and
inappropriate fire-fighting techniques are used to quench such a fire. As the temperature of burning
sodium chlorite is around 2200°C, water quickly turns to steam. Because thermal breakdown
products of sodium chlorite at high temperatures include molecular oxygen, appropriate techniques
are required to correctly extinguish closed containers or large amounts of dry material that has been
ignited.
Stratification of sodium chlorite in holding tanks may also occur and would influence the chlorine
dioxide yield. If stratification occurs in the bulk tank, sodium chlorite changes from high density to
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4. CHLORINE DIOXIDE
low density as it is fed. The density will continue to change until the material is re-mixed. In
stratified tanks, excess chlorite would be fed to the generator since the bottom of the tank will have
denser material, and this material would have more chlorite than required. Similarly, the bulk tank
would later discharge too little chlorite.
Although infrequent, such stratification is not readily apparent and may likely remain unnoticed by
operations unless the generator performance is evaluated frequently. If stratification or
crystallization occurs in bulk delivery trucks, the entire content should be warmed prior to delivery
so that the sodium chlorite is re-mixed. Operators should be aware of the possibility of stratification
and crystallization during delivery conditions.
Sodium chlorite is commercially available as a 38 percent or 25 percent solution. Chemical and
physical properties are given in Table 4-4.
Table 4-4. Properties of Sodium Chlorite as Commercially Available
Sodium Chlorite, (%) NaCI02
Sodium Chloride, (%) NaCI
Inert Ingredients, mixture of other sodium salts (%)
Water (%)
Appearance
Density @ 35°C (Ib/gal), typical
Crystallization Point (°C)
38% Solution*
38
1.5-7.5
3-4
55-61
Slightly cloudy, pale
yellow
11.4
25
25% Solution*
25
1-4.5
3-4
68-74
Clear, pale yellow
10.1
-7
" Source: Vulcan Chemicals
For systems handling the 38 percent solution, storage tanks, piping and pumps will require a heated
enclosure, or heat tracing and insulation. The 25 percent solution may not require any special
protection except in cold climates.
The ideal production of 1.0 pound of~chlorine dioxide requires 0.5 pounds of chlorine and 1.34
pounds of pure sodium chlorite. Chlorine gas is available as a nearly 100 percent pure chemical on a
weight basis. Gas flow metering devices are typically limited to +/- 5 percent accuracy at full rated
capacity. For example, a 100 pound per day flow tube would allow between 20 and 30 pounds of
chlorine to flow if set at 25 pounds per day (i.e., 25 +/- 5 percent of maximum flow capacity).
Sodium chlorite is supplied commercially as a pre-mixed aqueous solution of various strengths. The
25 percent solution is the most commonly used grade for potable water treatment.
Pure chlorine dioxide solutions (very dark amber and oily in appearance) are very dangerous and are
likely to detonate if exposed to oxidizable materials or vapors, or even to bright lights. They are
extremely uncommon except perhaps in very specific laboratory setup systems using concentrated
sodium chlorite and concentrated acid mixtures. Such laboratory generation methods are not
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4. CHLORINE DIOXIDE
recommended for the uninitiated laboratory analyst or operator. Inexperienced personnel should not
mix strong acid and strong sodium chlorite solutions together unless they are familiar with the
purgeable extraction methods for sodium chlorite and have a safely designed setup under a fume
hood.
4.8 Summary
4.8.1 Advantages and Disadvantages of Chlorine Dioxide Use
The following list highlights selected advantages and disadvantages of using chlorine dioxide as a
disinfection method for drinking water (Masschelein, 1992; DeMers and Renner, 1992, Gallagher et
al., 1994). Because of the wide variation of system size, water quality, and dosages applied, some of
these advantages and disadvantages may not apply to a particular system.
Advantages
Chlorine dioxide is more effective than chlorine and chloramines for inactivation of viruses,
Cryptosporidium, and Giardia.
Chlorine dioxide oxidizes iron, manganese, and sulfides.
Chlorine dioxide may enhance the clarification process.
Taste and odors resulting from algae and decaying vegetation, as well as phenolic compounds, are
controlled by chlorine dioxide.
Under proper generation conditions (i.e., no excess chlorine), halogen-substituted DBFs are not
formed.
Chlorine dioxide is easy to generate.
Biocidal properties are not influenced by pH.
Chlorine dioxide provides residuals.
Disadvantages
The chlorine dioxide process forms the specific byproducts chlorite and chlorate.
Generator efficiency and optimization difficulty can cause excess chlorine to be fed at the application
point, which can potentially form halogen-substitute DBFs.
Costs associated with training, sampling, and laboratory testing for chlorite and chlorate are high.
Equipment is typically rented, and the cost of the sodium chlorite is high.
Measuring chlorine dioxide gas is explosive, so it must be generated on-site.
Chlorine dioxide decomposes in sunlight.
Chlorine dioxide must be made on-site.
Can lead to production noxious odors in some systems.
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4.8.2 Summary Table
Table 4-5 summarizes considerations and descriptions for chlorine dioxide use.
Table 4-5. Summary for Chlorine Dioxide
Consideration
Inactivalion Efficiency
Byproducts Formation
Point of Application
Special Considerations
Description
Generation Chlorine dioxide must be generated on-site. In most potable water applications, chlorine
dioxide is generated as needed and directly educed or injected into a diluting stream.
Generators are available that utilize sodium chlorite and a variety of feedstocks such as
Cla gas, sodium hypochlorite, and sulfuric or hydrochloric acid. Small samples of
generated solutions, up to 1 percent (10 g/L) chlorine dioxide can be safely stored if the
solution is protected from light, chilled (<5°C), and has no unventilated headspace.
Primary Uses Chlorine dioxide is utilized as a primary or secondary disinfectant, for taste and odor
control, TTHM/HAA reduction, Fe and Mn control, color removal, sulfide and phenol
destruction, and Zebra mussel control.
Chlorine dioxide rapidly inactivates most microorganisms over a wide pH range. It is
more effective than chlorine (for pathogens other than viruses) and is not pH dependent
between pH 5-10, but is less effective than ozone.
When added to water, chlorine dioxide reacts with many organic and inorganic
compounds. The reactions produce chlorite and chlorate as endproducts (compounds
that are suspected of causing hemolytic anemia and other health effects). Chlorate ion
is formed predominantly in downstream reactions between residual chlorite and free
chlorine when used as the distribution system disinfectant. Chlorine dioxide does not
produce THMs. The use of chlorine dioxide aids in reducing the formation of TTHMs
and HAAs by oxidizing precursors, and by allowing the point of chlorination to be moved
farther downstream in the plant after coagulation, sedimentation, and filtration have
reduced the quantity of NOW.
In conventional treatment plants, chlorine dioxide used for oxidation is fed either in the
raw water, in the sedimentation basins, or following sedimentation. To limit the oxidant
demand, and therefore chlorine dioxide dose and the formation of chlorite, it is common
to add chlorine dioxide following sedimentation. Concerns about possible taste and odor
complaints have limited the use of-chlorine dioxide to provide a disinfectant residual in
the distribution system. Consequently, public water suppliers that are considering the
use of chlorine dioxide for oxidation and primary disinfectant applications may want to
consider chloramines for secondary disinfection.
An oxidant demand study should be completed to determine an approximate chlorine
dioxide dosage to obtain the required CT value as a disinfectant. In addition to the toxic
effects of chlorine, chlorine dioxide gas is explosive at levels > 10% in air. The chlorine
dioxide dosage cannot exceed 1.4 mg/L to limit the total combined concentration of
CI02, CI02-, CI03-, to a maximum of 1.0 mg/L. Under the proposed DBP regulations,
the MRDL for chlorine dioxide is 0.8 mg/L and the MCL for chlorite is 1.0 mg/L.
Regulations concerning the use of chlorine dioxide vary from state-to-state.
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4.9 References
1. Aieta, E., and J.D.Berg. 1986. j"A Review of Chlorine Dioxide in Drinking Water Treatment."
J. AWWA. 78(6):62-72.
2. Aieta, E.M., P.V. Roberts, and M. Hernandez. 1984. "Determination of Chlorine Dioxide,
Chlorine and Chlorate in Water." J. AWWA. 76(1):64-70.
3. Alvarez, M.E. and R.T. O'Brien. 1982. "Mechanism of Inactivation of Poliovirus by Chlorine
Dioxide and Iodine." Appl. Envir. Microbiol. 44:1064.
4. AWWA (American Water Works Association). 1995. AWWA Standard B303-95: Sodium
Chlorite.
5. AWWA. 1991. Guidance Manual for Compliance with the Filtration and Disinfection
Requirements for Public Water Systems Using Surface Water Sources.
6. AWWA. 1990. Water Quality and Treatment, fourth edition. McGraw-Hill, Inc., New York, NY.
7. Bedulivich, T.S., M.N. Svetlakova, and N.N. Trakhtman. 1954. "Use of Chlorine Dioxide in
Purification of Water." Chemical Abstracts. 48:2953.
8. Bernarde, M.A., et al. 1967a. "Kinetics and Mechanism of Bacterial Disinfection by Chlorine
Dioxide." /. Appl. Microbiol 15(2):257.
9. Bernarde, M.A., W.B. Snow, and V.P. Olivieri. 1967b. "Chlorine Dioxide Disinfection
Temperature Effects." J. Appl. Bacterial. 30(1): 159.
10. Chen, Y.S.R., O.J. Sproul, and A.J. Rubin. 1985. "Inactivation ofNaegleria gruberi Cysts by
Chlorine Dioxide." Water Res. 19(6):783.
11. Chen, Y.S.R., O.J. Sproul, and A.J. Rubin. 1984. "Inactivation ofNaegleria Gruberi cysts by
Chlorine Dioxide." EPA Grant R808150-02-0, Department of Civil Engineering, Ohio State
University.
12. CRC Handbook of Chemistry and Physics. 1990. D.L. Lide (editor), Seventy-first edition, CRC
Press, Boca Raton, FL.
13. Cronier, S., et al. 1978. Water Chlorination Environmental Impact and Health Effects, Vol. 2. R.
L. Jolley, et al. (editors) Ann Arbor Science Publishers, Inc. Ann Arbor, MI.
14. Demers, L.D., and R. Renner. 1992. Alternative Disinfectant Technologies for Small Drinking
Water Systems. AWWARF, Denver, CO.
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15. Dixon, K.L. and R.G. Lee. 1991. "Disinfection By-Products Control: A Survey of American
System Treatment Plants." Presented at AWWA Conference, Philadelphia, PA.
16. Emmenegger, F. and G. Gordon. 1967. "The Rapid Interaction Between Sodium Chlorite and
Dissolved Chlorine." Inorg. Chem. 6(3):633.
17. Finch, G.R., L.R. Liyanage, M. Belosevic, and L.L. Gyiirek. 1997. "Effects of Chlorine Dioxide
Preconditioning on Inactivation of Cryptosporidium by Free Chlorine and Monochloramine:
Process Design Requirements." Proceedings 1996 Water Quality Technology Conference; Part
II. Boston, MA.
18. Finch, G.R., L.R. Liyanage, and M. Belosevic. 1995. "Effect of Disinfectants and
Cryptosporidium and Giardia." Third International Symposium on Chlorine Dioxide: Drinking
Water, Process Water, and Wastewater Issues.
19. Gallagher, D.L., R.C. Hoehn, A.M. Dietrich. 1994. Sources, Occurrence, and Control of
Chlorine Dioxide By-Product Residuals in Drinking Water. AWWARF, Denver, CO.
20. Gates, D.J. 1998. The Chlorine Dioxide Handbook; Water Disinfection Series. AWWA
Publishing, Denver, CO.
21. Gates, DJ. 1989. "Chlorine Dioxide Generation Technology and Mythology." Conference
proceedings, Advances in Water Analysis and Treatment, AWWA, Philadelphia, PA.
22. Ghandbari, E. H., et al. 1983. "Reactions of Chlorine and Chlorine Dioxide with Free Fatty
Acids, Fatty Acid Esters, and Triglycerides." Water Chlorination: Environmental Impact and
Health Effects, R. L. Jolley, et al. (editors), Lewis, Chelsea, MI.
23. Gordon, G., G.L. Emmert, and B. Bubnis. 1995. "Bromate Ion Formation in Water When
Chlorine Dioxide is Photolyzed in the Presence of Bromide Ion." Conference proceedings,
AWWA Water Quality Technology Conference, New Orleans, LA.
24. Gordon, G., et al. 1990. "Minimizing Chlorite Ion and Chlorate Ion in Water Treated with
Chlorine Dioxide." J. AWWA. 82(4): 160-165.
25. Gordon, G., W.J. Cooper, R.G. Rice, and G.E. Pacey. 1987. Disinfectant Residual Measurement
Methods, AWWARF, Denver, CO.
26. Gordon, G., R.G. Kieffer, and D.H. Rosenblatt. 1972. "The Chemistry of Chlorine Dioxide."
Progress in Organic Chemistry, vol. 15. S.J. Lippaer (editor). Wiley Interscience, New York,
NY.
27. Great Lakes Upper Mississippi River Board of State Public Health (GLUMRB) and
Environmental Managers. 1992. Recommended Standards for Water Works, Health Research
Inc., Albany, NY.
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28. Gregory, D. and K. Carlson. 1998. "Applicability of Chlorine Dioxide for Cryptosporidium
Inactivation." Proceedings 1998 Water Quality Technology Conference, San Diego, CA.
29. Griese, M.H., K. Mauser, M. Berkemeier, and G. Gordon. 1991. "Using Reducing Agents to
Eliminate Chlorine Dioxide and Chlorite Ion Residuals in Drinking Water." /. AWWA. 83(5):56.
30. Hoehn, R.C. 1992. "Chlorine Dioxide Use in Water Treatment: Key Issues." Conference
proceedings, Chlorine Dioxide: Drinking Water Issues: Second International Symposium.
Houston, TX.
31. Hoehn, R.C., A.A. Rosenblatt, and D.J. Gates. 1996. "Considerations for Chlorine Dioxide
Treatment of Drinking Water." Conference proceedings, AWWA Water Quality Technology
Conference, Boston, MA.
32. Hofman, R., R.C. Andrews, and Q. Ye. 1997. "Chlorite Formation When Disinfecting Drinking
Water to Giardia Inactivation Requirements Using Chlorine Dioxide." Conference proceedings,
ASCE/CSCE Conference, Edmonton, Alberta, July.
33. Knocke, W.R. and A. latrou. 1993. Chlorite Ion Reduction by Ferrous Ion Addition. AWWARF,
Denver, CO.
34. Korich, D.G., et al. 1990. "Effects of Ozone, Chlorine Dioxide, Chlorine, and Monochloramine
on Cryptosporidium parvum oocyst Viability." Appl. Environ. Microbiol. 56:1423-1428.
35. LeChevallier, M.W., et al. 1997. "Chlorine Dioxide for Control of Cryptosporidium and
Disinfection Byproducts." Conference proceedings, 1996 AWWA Water Quality Technology
Conference Part II, Boston, Massachusetts.
36. LeChevallier, M.W., et al. 1996. "Chlorine Dioxide for Control of Cryptosporidium and
Disinfection Byproducts." Conference proceedings, AWWA Water Quality Technology
Conference, Boston, Massachusetts.
37. Liyanage, L.R.J, et al. 1997. "Effects of Aqueous Chlorine and Oxychlorine Compounds on
Cryptosporidium Parvum Oocysts." Environ. Sci. & Tech. 31(7): 1992-1994
38. Masschelein, W.J. 1992. "Unit Processes in Drinking Water Treatment." Marcel Decker D.C.,
New York, Brussels, Hong Kong.
39. Merkle, J.C. and C.B. Reeverts. 1997. "Ground Water Treatment: What Are the States Doing
Now?" AWWARF, Denver, CO.
40. Noack, M.G. and R.L. Doerr. 1977. "Reactions of Chlorine, Chlorine Dioxide and Mixtures of
Humic Acid: An Interim Report." Conference proceedings, Second Conference on the
Environmental Impact of Water Chlorination. R.L. Jolley, H. Gorchev, and D. Hey ward (editors),
Gatlinburg, TN.
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41. Noss, C.I., W.H. Dennis, V.P. Olivieri. 1983. "Reactivity of Chlorine Dioxide with Nucleic
Acids and Proteins." Water Chlorination: Environmental Impact and Health Effects. R. L. Jolley,,
et al. (editors), Lewis Publishers, Chelsea, MI.
42. Olivieri, V.P., et al. 1985. "Mode of Action of Chlorine Dioxide on Selected Viruses." Water
Chlorination: Environmental Impact and Health Effects. R. L. Jolley, et al. (editors), Lewis,
Chelsea, MI.
43. Olivieri, V.P., et al. 1984. Stability and Effectiveness of Chlorine Disinfectants in Water
Distribution Systems. USEPA, Cincinnati, OH.
44. Peeters, J. E. et al. 1989. "Effect of Disinfection of Drinking Water with Ozone or Chlorine
Dioxide on Survival of Cryptosporidiumparvum oocysts." Appl. Environ. Microbiol. r5:1519-
1522.
45. Pitochelli, A. 1995. "Chlorine Dioxide Generation Chemistry." Conference proceedings, Third
International Symposium, Chlorine Dioxide: Drinking Water, Process Water, and Wastewater
Issues. New Orleans, LA.
46. Ransome, M.E., T.N. Whitmore, and E.G. Carrington. 1993. "Effect of Disinfectants on the
Viability of Cryptosporidium parvum Oocysts." Water Supply. 11(1):103-117.
47. Rav-Acha, C., A. Serri, E. Choshen, B. Limoni. 1984. "Disinfection of Drinking Water Rich in
Bromide with Chlorine and Chlorine Dioxide, While Minimizing the Formation of Undesirable
Byproducts." Wat. Sci. Technol. 17:611.
48. Richardson, S.D. et al. 1994. "Multispectral Identification of C1O2 Disinfection Byproducts in
Drinking Water." Environ. Sci. & Technol. 28(4):592-599.
49. Ridenour, G.M. and E.H. Armbruster. 1949. "Bactericidal Effects of Chlorine Dioxide."
J. AWWA. 41:537.
50. Ridenour, G. M. and R.S. Ingols. 1947. "Bactericidal Properties of Chlorine Dioxide."
/. AWWA. 39.
51. Ridenour, G.M., and R.S. Ingols. 1946. "Inactivation of Poliomyelitis Virus by Free Chlorine."
Amer. Public Health. 36:639.
52. Ridenour, G.M., and R.S. Ingols, and E.H. Armbruster. 1949. "Sporicidal Properties of Chlorine
Dioxide." Water & Sewage Works. 96(8):279.
53. Roberts, P.V., E.M. Aieta, J.D. Berg, and B.M. Chow. 1980. "Chlorine Dioxide for Wastewater
Disinfection: A Feasibility Evaluation." Stanford University Technical Report 251. October.
EPA Guidance Manual 4-38 April 1999
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4. CHLORINE DIOXIDE
54. Roller, S. D. et al. 1980. "Mode of Bacterial Inactivation by Chlorine Dioxide." Water Res.
14:635.
55. Singer, P.C. 1992. "Formation and Characterization of Disinfection Byproducts." Presented at
the First International Conference on the Safety of Water Disinfection: Balancing Chemical and
Microbial Risks.
56. Singer, P.C., and W.K. O'Neil. 1987. "Technical Note: The Formation of Chlorate from the
Reaction of Chlorine and Chlorite in Dilute Aqueous Solution." /. A WWA. 79(11):75.
57. Smith, J. E., and J.L. McVey. 1973. "Virus Inactivation by Chlorine Dioxide and Its Application
to Storm Water Overflow." Proceeding, ACS annual meeting. 13(2): 177.
58. Sobsey, M. 1988. "Detection and Chlorine Disinfection of Hepatitis A in Water." CR-813-024,
EPA Quarterly Report, December.
59. Sproul, O. J. et al. 1983. "Comparison of Chlorine and Chlorine Dioxide for Inactivation of
Amoebic Cyst." Envir. Technol. Letters. 4:335.
60. Thompson, A.L. 1989. "Practical Considerations for Application of Chlorine Dioxide in
Municipal Water Systems." Conference proceedings,, Chlorine Dioxide Workshop. AWWARF,
CMA, EPA. Denver, CO.
61. Trakhtman, N.N. 1949. "Chlorine Dioxide in Water Disinfection." Chemical Abstracts. 43:1508.
62. USEPA (U.S. Environmental Protection Agency). 1983. "Trihalomethanes in Drinking Water:
Sampling, Analysis, Monitoring, and Compliance." EPA 570/9-83-002, August.
63. USEPA. 1979. "Effect of Particulates on Disinfection of Enteroviruses and Coliform Bacteria in
Water by Chlorine Dioxide." EPA-600/2-79-054.
64. USEPA 1978. "Effect of Particulates on Inactivation of Enteroviruses in Water by Chlorine
Dioxide." EPA-600/9-79-018, Cincinnati, OH.
65! Werdehoff, K.S, and P.C. Singer. 1987. "Chlorine Dioxide Effects on THMFP, TOXFP and the
Formation of Inorganic By-Products." J. AWWA. 79(9): 107.
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5. POTASSIUM PERMANGANATE
Potassium permanganate (KMnO4) is used primarily to control taste and odors, remove color, control
biological growth in treatment plants, and remove iron and manganese. In a secondary role,
potassium permanganate may be useful in controlling the formation of THMs and other DBFs by
oxidizing precursors and reducing the demand for other disinfectants (Hazen and Sawyer, 1992).
The mechanism of reduced DBFs may be as simple as moving the point of chlorine application
further downstream in the treatment train using potassium permanganate to control taste and odors,
color, algae, etc. instead of chlorine. Although potassium permanganate has many potential uses as
an oxidant, it is a poor disinfectant.
5.1 Potassium Permanganate Chemistry
5.1.1 Oxidation Potential
Potassium permanganate is highly reactive under conditions found in the water industry. It will
oxidize a wide variety of inorganic and organic substances. Potassium permanganate (Mn 7+) is
reduced to manganese dioxide (MnO2) (Mn 4+) which precipitates out of solution (Hazen and
Sawyer, 1992). All reactions are exothermic. Under acidic conditions the oxidation half-reactions
are (CRC, 1990):
MnO4- + 4/T + 3e -> MnO2 + 2H2O E" = 1.68V
MnOi + 8/T + 5e^ Mn2+ + 4H2O £"=1.5 IV
Under alkaline conditions, the half-reaction is (CRC, 1990):
MnO4- + 2H2O + 3e -> MnO2 + 4OH' E" = 0.60V
Reaction rates for the oxidation of constituents found in natural waters are relatively fast and depend
on temperature, pH, and dosage.
5.1.2 Ability To Form a Residual
It is not desirable to maintain a residual of KMnO4 because of its tendency to give water a pink color.
5.2 Generation
Potassium permanganate is only supplied in dry form. A concentrated KMnO4 solution (typically 1
to 4 percent) is generated on-site for water treatment applications; the solution is pink or purple in
color. KMnO4 has a bulk density of approximately 100 lb/ft3 and its solubility in water is 6.4 g/mL
at 20°C.
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Depending on the amount of permanganate required, these solutions can be made up in batch modes,
using dissolver/storage tanks with mixers and a metering pump for small feed systems. Larger
systems will include a dry chemical feeder, storage hopper and dust collector configured to
automatically supply permanganate to the solution dissolver/storage tank.
KMnC>4 solution is made up of dry crystalline permanganate solids added to make-up water and then
stirred to obtain the desired permanganate concentration. The cost of KMnCU ranges from $1.50 to
$2.00 per pound (1997 costs), depending on the quantity ordered. Shipment containers are typically
buckets or drums. Potassium permanganate is supplied in various grades. Pure KMnC>4 is non-
hygroscopic but technical grades will absorb some moisture and will have a tendency to cake
together. For systems using dry chemical feeders, a free-flowing grade is available that contains anti-
caking additives (Hazen and Sawyer, 1992).
Potassium permanganate is a strong oxidizer and should be carefully handled when preparing the
feed solution. No byproducts are generated from making the solution. However, this dark
purple/black crystalline solid can cause serious eye injury, is a skin and inhalation irritant, and can be
fatal if swallowed. As such, special handling procedures include the use of safety goggles and a face
shield, an MSA™/NIOSH approved dust mask, and wearing impervious gloves, coveralls, and boots
to minimize skin contact.
5.3 Primary Uses and Points of Application
Although potassium permanganate can inactivate various bacteria and viruses, it is not used as a
primary or secondary disinfectant when applied at commonly used treatment levels. Potassium
permanganate levels that may be required to obtain primary or secondary disinfection could be cost
prohibitive. However, potassium permanganate is used in drinking water treatment to achieve a
variety of other purposes including:
• Oxidation of iron and manganese;
• Oxidation of taste and odor compound;
• Control of nuisance organisms; and
• Control of DBP formation.
5.3.1 Primary Uses
5.3.1.1 Iron and Manganese Oxidation
A primary use of permanganate is iron and manganese removal. Permanganate will oxidize iron and
manganese to convert ferrous (2+) iron into the ferric (3+) state and 2+ manganese to the 4+ state.
The oxidized forms will precipitate as ferric hydroxide and manganese hydroxide (AWWA, 1991).
The precise chemical composition of the precipitate will depend on the nature of the water,
temperature, and pH.
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The classic reactions for the oxidation of iron and manganese are:
3Fe2+ + KMnO4 + 1H2O •* 3Fe(OH)3M + MnO2(s, + K* + 5H+
3Mn2+ + 2KMnO4 + 2H2O •* 5MnO2(s) + 2K+ + 4H+
These reactions show that alkalinity is consumed through acid production at the rate of 1.49 mg/L as
CaCO3 per mg/L of Fe+2 and 1.21 mg/L as CaCO3 per mg/L of Mn"1"2 oxidized. This consumption of
alkalinity should be considered when permanganate treatment is used along with alum coagulation,
which also requires alkalinity to form precipitates.
The potassium permanganate dose required for oxidation is 0.94 mg/mg iron and 1.92 mg/mg
manganese (Culp/Wesner/Culp, 1986). In practice, the actual amount of potassium permanganate
used has been found to be less than that indicated by stoichiometry. It is thought that this is because
of the catalytic influence of MnO2on the reactions (O'Connell, 1978). The oxidation time ranges
from 5 to 10 minutes, provided that the pH is over 7.0 (Kawamura, 1991).
5.3.1.2 Oxidation of Taste and Odor Compounds
Potassium permanganate is used to remove taste and odor causing compounds. Lalezary et al. (1986)
used permanganate to treat earthy-musty smelling compounds in drinking water. Doses of potassium
permanganate used to treat taste and odor causing compounds range from 0.25 to 20 mg/L.
5.3.1.3 Control of Nuisance Organisms
Asiatic Clams
Cameron et al. (1989) investigated the effectiveness of potassium permanganate to control the
Asiatic clam in both the juvenile and adult phases. The adult Asiatic clam was found to be much
more resistant to permanganate than the juvenile form. Potassium permanganate doses used to
control the juvenile Asiatic clam range from 1.1 to 4.8 mg/L.
Zebra Mussels
Klerks and Fraleigh (1991) evaluated the effectiveness of permanganate against adult zebra mussels.
Continuous potassium permanganate dosing of 0.5 to 2.5 mg/L proved to be the most effective.
5.3.7.4 DBP Control
It is anticipated that potassium permanganate may play a role in disinfection and DBP control
strategies in water treatment. Potassium permanganate could be used to oxidize organic precursors at
the head of the treatment plant minimizing the formation of byproducts at the downstream
disinfection stage of the plant (Hazen and Sawyer, 1992). Test results from a study conducted at two
water treatment plants in North Carolina (Section 5.5.1) showed that pretreatment with permanganate
reduced chloroform formation; however, the reduction was small at doses typically used at water
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treatment plants. The study also indicated that pre-oxidation with permanganate had no net effect on
the chlorine demand of the water (Singer et al., 1980).
5.3.2 Points of Application
In conventional treatment plants, potassium permanganate solution is added to the raw water intake,
at the rapid mix tank in conjunction with coagulants, or at clarifiers upstream of filters. In direct
filtration plants, this oxidant is typically added at the raw water intake to increase the contact time
upstream of the filter units (Montgomery, 1985). In all cases, potassium permanganate is added prior
to filtration.
Potassium permanganate solution is typically pumped from the concentrated solution tank to the
injection point. If the injection point is a pipeline, a standard injection nozzle protruding midway
into the pipe section is used. Injection nozzles can also be used to supply the solution to mixing
chambers and clarifiers. Permanganate is a reactive, fast-acting oxidizer and does not require special
mixing equipment at the point of injection to be effective.
5.3.2.1 Impact on Other Treatment Processes
The use of potassium permanganate has little impact on other treatment processes at the water
treatment facility. See Section 5.7 for permanganate operational considerations.
5.4 Pathogen Inactivation and Disinfection Efficacy
Potassium permanganate is an oxidizing agent widely used throughout the water industry. While it is
not considered a primary disinfectant, potassium permanganate has an effect on the development of a
disinfection strategy by serving as an alternative to pre-chlorination or other oxidants at locations in a
treatment plant where chemical oxidation is desired for control of color, taste and odor, and algae.
5.4.1 Inactivation Mechanisms
The primary mode of pathogen inactivation by potassium permanganate is direct oxidation of cell
material or specific enzyme destruction (Webber and Posselt, 1972). In the same fashion, the
permanganate ion (MnCV) attacks a wide range of microorganisms such as bacteria, fungi, viruses,
and algae.
Application of potassium permanganate results in the precipitation of manganese dioxide. This
mechanism represents an additional method for the removal of microorganisms from potable water
(Cleasby et al., 1964). In colloidal form, the manganese dioxide precipitant has an outer layer of
exposed OH groups. These groups are capable of adsorbing charged species and particles in addition
to neutral molecules (Posselt et al., 1967). As the precipitant is formed, microorganisms can be
adsorbed into the colloids and settled.
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5.4.2 Environmental Effects
Inactivation efficiency depends upon the permanganate concentration, contact time, temperature, pH,
and presence of other oxidizable material. Several of the key parameters are discussed below.
5.4.2.1 pH
Alkaline conditions enhance the capability of potassium permanganate to oxidize organic matter;
however, the opposite is true for its disinfecting power. Typically, potassium permanganate is a
better biocide under acidic conditions than under alkaline conditions (Cleasby et al., 1964 and
Wagner, 1951). Results from a study conducted in 1964 indicated that permanganate generally was a
more effective biocide for E. coli at lower pHs, exhibiting more than a 2-log removal at a pH of 5.9
and a water temperature of both 0 and 20°C (Cleasby et al., 1964). In fact, Cleasby found that pH is
the major factor affecting disinfection effectiveness with potassium permanganate. As such, natural
waters with pH values of 5.9 or less would be conducive to potassium permanganate disinfection,
particularly as a substitute for prechlorination. Moreover a study conducted at the University of
Arizona found that potassium permanganate will inactivate Legionella pneumophila more rapidly at
pH 6.0 than at pH 8.0 (Yahya et al., 1990a).
These results are consistent with earlier results concerning the effects of pH on commercial antiseptic
performance (Hazen and Sawyer, 1992). In general, based on the limited results from these studies,
disinfection effectiveness of potassium permanganate increases with decreasing pH.
5.4.2.2 Temperature
Higher temperatures slightly enhance bactericidal action of potassium permanganate. The results
from a study conducted on polio virus showed that oxidation deactivation is enhanced by higher
temperatures (Lund, 1963). These results are consistent with results obtained for E. coli. inactivation
(Cleasby et al., 1964).
5.4.2.3 Dissolved Or games and Inorganics
The presence of oxidizable organics or inorganics in the water reduces the disinfection effectiveness
of this disinfectant because some of the applied potassium permanganate will be consumed in the
oxidation of organics and inorganics. Permanganate oxidizes a wide variety of inorganic and organic
substances in the pH range of 4 to 9. Under typical water conditions, iron and manganese are
oxidized and precipitated and most contaminants that cause odors and tastes, such as phenols and
algae, are readily degraded by permanganate (Hazen and Sawyer, 1992).
5.4.3 Use as a Disinfectant
A number of investigations have been performed to determine the relative capability of potassium
permanganate as a disinfectant. The following sections contain a description of the disinfection
efficiency of potassium permanganate in regards to bacteria, virus, and protozoa inactivation.
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5.4.3.1 Bacteria Inactivation
High dosage rates were required to accomplish complete inactivation of bacteria in three studies.
Early research showed that a dose of 2.5 mg/L was required for complete inactivation of coliform
bacteria (Le Strat, 1944). In this study, water from the Marne River was dosed with potassium
permanganate at concentrations of 0 to 2.5 mg/L. Following mixing, the samples were placed in a
darkened room for 2 hours at a constant temperature of 19.8°C.
Banerjea (1950) investigated the disinfectant ability of potassium permanganate on several
waterborne pathogenic microorganisms. The investigation studied Vibrio cholerae, Salm. typhi, and
Bact. flexner. The results indicated that doses of 20 mg/L and contact times of 24 hours were
necesSary to deactivate these pathogens; however, even under these conditions the complete absence
of Salm. typhi or Bact. flexner was not assured, even at a potassium permanganate concentration that
turned the water an objectionable pink color.
Results from a study conducted in 1976 at the Las Vegas Valley Water District/Southern Nevada
System of Lake Mead water showed that complete removal of coliform bacteria were accomplished
at doses of 1, 2, 3,4, 5, and 6 mg/L (Hazen and Sawyer, 1992). Contact times of 30 minutes were
provided with doses of 1 and 2 mg/L, and 10 minutes contact times were provided for higher dosages
in this study.
5.4.3.2 Virus Inactivation
Potassium permanganate has been proven effective against certain viruses. A dose of 50 mg/L of
potassium permanganate and a contact time of 2 hours was required for inactivation of poliovirus
(strain MVA) (Hazen and Sawyer, 1992). A "potassium" permanganate dose of 5.0 mg/L and a
contact time of 33 minutes was needed for 1-log inactivation of type 1 poliovirus (Yahya et al.,
1990b). Tests showed a significantly higher inactivation rate at 23°C than at 7°C; however, there was
no significant difference in activation rates at pH 6.0 and pH 8.0.
Potassium permanganate doses from 0.5 to 5 mg/L were capable of obtaining at least a 2 log
inactivation of the surrogate virus, MS-2 bacteriophage with E. coli as the host bacterium (Yahya et
al., 1989). Results showed that at pH 6.0 and 8.0, a 2-log inactivation occurred after a contact time
of at least 52 minutes and a residual of 0.5 mg/L. At a residual of 5.0 mg/L, approximately 7 and 13
minutes were required for 2-log inactivation at pHs of 8.0 and 6.0, respectively. These results
contradict the previously cited studies that potassium permanganate becomes more effective as the
pH decreases.
5.4.3.3 Protozoa Inactivation
No information pertaining to protozoa inactivation by potassium permanganate is available in the
literature. However, based on the other disinfectants discussed in this report, protozoa are
significantly more resistant than viruses; therefore, it is likely that the dosages and contact times
required for protozoa inactivation would be impractical.
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5.4.3.4 CT Curves
Table 5-1 shows CT values for the inactivation of bacteriophage MS-2. These data have been
provided as an indication of the potential of potassium permanganate. These values are somewhat
inconsistent and do not include a safety factor and should not be used to establish CT requirements.
Table 5-1. Potassium Permanganate CT Values for 2-log Inactivation of
MS-2 Bacteriophage
Residual
(mg/L)
0.5
1.5
2
5
pH 6.01
(mg min / L)
27.4 (a)
32.0 (a)
63.8 (a)
pH8.01
(mg min / L)
26.1 (a)
50.9 (b)
53.5 (c)
35.5 (c)
Source: USEPA, 1990.
Note: ' Letters indicate different experimental conditions.
A 1990 study investigated CT values for Legionella pneumophila inactivation. CT values for 99
percent (2-log) inactivation of Legionella pneumophila at pH 6.0 were determined to be 42.7 mg
min/L at a dose of 1.0 mg/L (contact time 42.7 minutes) and 41.0 mg min/L at a dose of 5.0 mg/L
(contact time 8.2 minutes) (Yahya et al., 1990a).
5.5 Disinfection Byproduct Formation
No literature is available that specifically addressed DBFs when using potassium permanganate.
However, several studies have, been conducted with water treatment plants that have replaced the pre-
chlorination process with potassium permanganate and relocated the point of chlorine addition for
post-treatment disinfection. Pretreatment with permanganate in combination with post-treatment
chlorination will typically result in lower DBF concentrations than would otherwise occur from
traditional pre-chlorination (Ficek and Boll, 1980; and Singer et al., 1980). Under this approach,
potassium permanganate serves as a substitute for chlorine to achieve oxidation and may also reduce
the concentration of natural organic matter (NOM). However, systems should evaluate the impact on
CT values before moving the point of chlorination. The following subsections summarize the
outcomes of two studies.
5.5.1 Chapel-Hill and Durham, North Carolina Water Treatment
Plants
An investigation was conducted at the Chapel-Hill and Durham Water Treatment Plants to evaluate
the effects of potassium permanganate pretreatment on trihalomethane formation (Singer et al.,
1980). The Chapel-Hill Water Treatment Plant uses pre-chlorination prior to the rapid mix tank. At
the Durham Water Treatment Plant, chlorine is not added until after the sedimentation basin prior to
the filtration. Both are surface water treatment plants, treating water with low concentrations of
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alkalinity. Both sources of water are known to have high trihalomethane formation potentials
(Young and Singer, 1979).
Raw water samples taken from Chapel-Hill were found to contain relatively high turbidities, ranging
from 46 to 110 NTU and total organic carbon (TOC) concentrations ranging from 5.6 to 8.9 mg/L.
The Durham samples were coagulated then allowed to settle, which resulted in better water quality
than the Chapel-Hill samples. Following settling, this sample had a turbidity of 6.4 NTU and a TOC
of 2.9 mg/L. Sulfuric acid and sodium hydroxide were used to adjust the sample pH to either 6.5 or
10.3. These pH values were selected because they encompass the pH range typically found in
surface water coagulation-filtration and lime-softening treatment plants.
Potassium permanganate doses of 2 and 5 mg/L were found to be totally consumed within 1 and 4
hours, respectively, by the Chapel-Hill samples. At doses of 2 and 5 mg/L, the potassium
permanganate demand of the Durham samples after 4 hours were approximately 1.3 and 1.8 mg/L,
respectively.
. tn
This difference in permanganate demands between the Chapel-Hill and Durham samples may be
attributed to the water quality of the samples, in particular the TOC concentrations. TOC
measurements before and after the application of permanganate were approximately equal; however,
it is likely that the TOC after disinfection was at a higher oxidation state. Results of this study also
showed that permanganate is more reactive as an oxidant at higher pH values.
Despite the high degree of permanganate consumption, the reaction of permanganate appears to have
relatively little effect on chlorine demands. For example, consumption of 6 mg/L of permanganate
resulted in a chlorine demand reduction of approximately 1 mg/L. This observation suggests that
permanganate reacts with water impurities in a different manner, or at different sites, than chlorine.
One other possible explanation is that permanganate oxidizes certain organic substances, thereby
eliminating their chlorine demand and only partially oxidizing other organic substances making them
more reactive to chlorine.
Both the Chapel-Hill and Durham samples were tested for their chloroform formation potential. This
measurement is based on the amount of chloroform produced after seven days. The potential of the
Durham sample was reduced by 30 and 40 percent at pH 6.5 and 10.3, respectively, as a result of the
application of 10 mg/L of potassium permanganate for a period of 2 hours. Similar results were
obtained for the Chapel-Hill samples; however, the results at pH 6.5 did not show a reduction in
chloroform formation potential at low doses.
Two experiments were conducted on Chapel-Hill raw water to further explore the effects of low
doses of permanganate. The results indicated that permanganate has no effect on chloroform
production at doses up to 1 mg/L. At higher doses, chloroform formation potentials were reduced.
In summary, the key results obtained from the studies conducted at the Chapel-Hill and Durham
Water treatment plants were:
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• The reactivity of permanganate is a function of pH, permanganate dose, and raw water
quality.
• Permanganate reduces chloroform formation potentials. The reduction in the chloroform
formation potential is proportional to the amount of permanganate available after the initial
demand is overcome. Doses up to 1 mg/L were found to have no effect on chloroform
formation potentials.
• At pretreatment doses typically employed at water treatment plants, the effect of
permanganate on the overall chloroform production is relatively small. If permanganate is to
be used specifically to reduce trihalomethane formation, larger doses will be required.
However, one advantage for using permanganate for pretreatment is that the point of
application of chlorine can be shifted downstream of the sedimentation basins. This is likely
to result in fewer trihalomethane compounds.
5.5.2 American Water Works Association Research Foundation
TTHM Study
Another investigation examined the impacts of potassium permanganate addition on byproduct
formation at four water treatment plants (Ficek and Boll, 1980). All were conventional plants using
pre-chlorination in the treatment process. Plant design capacities ranged from 4.5 to 15 mgd.
Process modifications were made at each plant to replace the pre-chlorination facilities with
oxidation facilities for potassium permanganate addition. After the modifications were complete, an
AWWARF research team conducted a study to determine the impact of potassium permanganate
addition on total trihalomethane (TTHM) concentrations (George et al., 1990).
Prior to switching from pre-chlorination to pre-oxidation with potassium permanganate, average
daily TTHM concentrations at all four plants were between 79 and 99 jug/L. The average TTHM
concentration for all four plants was 92 ug/L. Following the conversion to potassium permanganate,
three of the four plants experienced greater than 30 percent reduction in TTHM concentrations. In
addition to TTHM reduction, potassium permanganate was found to oxidize taste and odor causing
compounds, iron and manganese, organic and inorganic matter, and reduce algal growth. Results
from the study also showed that the simultaneous application of potassium permanganate and
chlorine can increase THM formation.
5.6 Status of Analytical Methods
The atomic adsorption spectrophotometry method for the measurement of manganese is the preferred
method for measuring permanganate concentrations. Two colorimetric methods, persulfate and
periodate are also available (Standard Methods, 1995).
5.7 Operational Considerations
In utilizing potassium permanganate in water treatment, caution should be taken to prevent
overdosing, in which case, excess manganese will pass through the treatment plant. Proper dosing
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5. POTASSIUM PERMANGANATE
should be maintained to ensure that all of the permanganate is reduced (i.e., forming MnO2 solids)
and removed from the plant upstream of, or within, the filters. If residual manganese is reduced
downstream of the filters, the resulting solids can turn the finished water a brown/black color and
precipitate in the homes of consumers on heat exchange surfaces such as hot water heaters and
dishwashers.
Use of potassium permanganate can also be a source of manganese in the finished water, which is
regulated in drinking water with a secondary maximum contaminant level of 0.05 mg/L. Under
reducing conditions, the MnOa solids accumulated in filter backwash water and settling basins can be
reduced to soluble Mn2* and pass through the filters thereby remaining in the finished water.
Also, under these conditions, soluble Mn2+ in return water from settling basin dewatering facilities
and filter backwash water recycled to the head of the plant are potential sources of manganese that
will have to be treated and/or controlled to minimize finished water manganese levels (Singer, 1991).
Overdosing of permanganate in conventional plants is generally corrected by settling the excess
MnOi solids in the settling basin. Removal of the excess permanganate can be monitored
qualitatively by observing the disappearance of the pink color characteristic of permanganate. In
plants that do not utilize flocculation and sedimentation processes permanganate dosing should be
closely monitored (Montgomery, 1985).
In general, potassium permanganate does not interfere with other treatment processes or plant
conditions. Permanganate can be added downstream of, or concurrently with, coagulant and filter
polymer aids. Powdered activated carbon (PAC) and permanganate should not be added
concurrently. PAC should be added downstream of permanganate because it may consume
permanganate, rendering it unavailable for the oxidation of target organics. (Montgomery, 1985).
The space requirements for permanganate feed equipment vary depending on the type and size of
feed system. Dry feed systems require about half the floor area of batch systems because batch
systems typically have two dissolving tanks for redundancy. However, the head space requirements
are greater for dry feed systems where the storage hopper and dust collector are stacked on top of the
dry feeder (Kawamura, 1991). On-site storage of potassium permanganate also warrants some
consideration. Per OSHA requirements, oxidants such as permanganate should be stored separate
from organic chemicals such as polymers and activated carbon.
5.8 Summary
5.8,1 Advantages and Disadvantages of Potassium
Permanganate Use
The following list highlights selected advantages and disadvantages of using potassium
permanganate as a disinfection method for drinking water. Because of the wide variation of system
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size, water quality, and dosages applied, some of these advantages and disadvantages may not apply
to a particular system.
Advantages
• Potassium permanganate oxidizes iron and manganese.
• Potassium permanganate oxidizes odor and taste-causing compounds.
• Potassium permanganate is easy to transport, store, and apply.
• Potassium permanganate is useful in controlling the formation of THMs and other DBFs.
• Potassium permanganate controls nuisance organisms.
• The use of potassium permanganate has little impact on other treatment processes at the water
treatment facility.
• Potassium permanganate has been proven effective against certain viruses.
Disadvantages
• Long contact time is required.
• Potassium permanganate has a tendency to give water a pink color.
• Potassium permanganate is toxic and irritating to skin and mucous membranes.
• No byproducts are generated when preparing the feed solution, however this dark purple/black
crystalline solid can cause serious eye injury, is a skin and inhalation irritant, and can be fatal if
swallowed. Over-dosing is dangerous and may cause health problems such as chemical jaundice
and drop in blood pressure.
5.8.2 Summary Table
More research is needed regarding the disinfection properties and oxidation byproducts of
permanganate in water treatment. Also, a CT credit needs to be assigned to permanganate if it is to
be utilized as a disinfectant. However, given that alternative oxidants, such as ozone and chlorine
dioxide, demonstrate much greater efficacy in microbial control, permanganate is not likely to be
utilized as a primary oxidant for precursor control. Table 5-2 summarizes the information presented
in this chapter regarding the use of potassium permanganate in the drinking water treatment process.
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5. POTASSIUM PERMANGANATE
Table 5-2. Summary of Potassium Permanganate Use
Consideration
Description
Generation
Primary uses
Inactivation efficiency
Byproduct formation
Limitations
Points of application
Special considerations
Product supplied in dry form in buckets, drums, and bulk. On-site generation of solution is
required using chemical mixing and feed equipment.
Control of odor and taste, remove color, control biological growth, and remove iron and
manganese.
Not a good disinfectant. Can serve better as an alternative to chlorine or other disinfectants
where chemical oxidation is desired.
No literature was found that specifically addressed DBP formation from potassium
permanganate oxidation. Pretreatment with permanganate in combination with post-treatment
chlorination will typically result in lower DBP concentrations than would otherwise occur from
traditional pre-chlorination.
Not a good disinfectant; primarily used for pretreatment to minimize chlorine usage and
byproduct formation.
Conventional Treatment: raw water addition, rapid mix tank in conjunction with coagulants,
clarifiers upstream of filters. Direct Filtration: raw water intake. In all cases permanganate
should be added upstream of filters.
Caution should be taken to prevent overdosing. More research is needed to determine
disinfection properties and oxidation byproducts.
5.9 References
1. AWWA (American Water Works Association). 1991. Guidance Manual for Compliance with the
Filtration and Disinfection Requirements for Public Water Systems Using Surface^ Water
Sources.
2. Banerjea, R. 1950. "The Use of Potassium Permanganate in the Disinfection of Water." Ind.
Med. Gaz. 85:214-219.
3. Cameron, G.N., J.M. Symons, S.R. Spencer, and J.Y. Ma. 1989. "Minimizing THM Formation
During Control of the Asiatic Clam: A Comparison of Biocides." J. AWWA. 81(10):53-62.
4. Cleasby, J.L., E.R. Baumann, and C.D. Black. 1964. "Effectiveness of Potassium Permanganate
for Disinfection." J. AWWA. 56:466-474.
5. CRC. 1990. Handbook of Chemistry and Physics, seventy-first edition. D.L. Lide (editor). CRC
Press, Boca Raton, FL.
6. Culp/Wesner/Culp. 1986. Handbook of Public Water Systems. Van Nostrand Reinhold, New
York, NY.
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5. POTASSIUM PERMANGANATE
7. Ficek, K.J., and J.E. Boll. 1980. "Potassium Permanganate: An Alternative to Prechlorination."
Ague. 7:153-156.
8. George, D.B., V.D. Adams, S.A. Huddleston, K.L. Roberts, and M.B. Borup. 1990. Case Studies
of Modified Disinfection Practices for Trihalomethane Control, Potassium Permanganate.
AWWAR and AWWA, Denver, CO.
9. Hazen and Sawyer. 1992. Disinfection Alternatives for Safe Drinking Water. Van Nostrand
Reinhold, New York, NY.
10. Kawamura, S. 1991. Integrated Design of Water Treatment Facilities. John Wiley & Sons, Inc.,
New York, NY.
11. Klerks, P.L. and P.C. Fraleigh. 1991. "Controlling Adult Zebra Mussels with Oxidants."
J.AWWA. 83(12):92-100.
12. Lalezary, S., M. Pirbazari, and M.J. McGuire. 1986. "Oxidation of Five Earthy-Musty Taste and
Odor Compounds." J. AWWA. 78(3):62.
13. Le Strat. 1944. "Comparison des pouvoirs sterilisants du permanganate de potasses et de 1'eau de
javel a 1'egard d'eaux contaminees." Ann. Hygiene.
14. Lund, E. 1963. "Significance of Oxidation in Chemical Interaction of Polioviruses." Arch. Ges.
Virusdorsch. 12(5):648-660.
15. Montgomery, J.M. 1985. Water Treatment Principles and Design. John Wiley & Sons, Inc.,
New York, NY.
16. O'Connell, R.T. 1978. "Suspended Solids Removal." Water Treatment Plant Design. R.L. Sanks
(editor). Ann Arbor Science Publishers, Inc, Ann Arbor, MI.
17. Posselt, H.S., F. J. Anderson, and WJ. Webber. 1967^ "The Surface Chemistry of Hydrous
Manganese Dioxide." Presented at meeting of Water, Air, and Waste Chemistry Division,
American Chemical Society, Bar Harbor, FL, April.
18. Singer, P.C. 1991. "Research Needs for Alternative Oxidants and Disinfectants." Presented at the
Annual AWWA Conference, Philadelphia, June 23-27.
19. Singer, P.C., J.H. Borchardt, and J.M. Colthurst. 1980. "The Effects of Permanganate
Pretreatment on Trihalomethane Formation in Drinking Water." /. AWWA. 72(10):573-578.
20. Standard Methods. 1995. Standard Methods for the Examination of Water and Wastewater,
nineteenth edition. American Public Health Association, AWWA, and Water Pollution Control
Fed., Washington, D.C.
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5. POTASSIUM PERMANGANATE
21. USEPA. 1990. Guidance Manual for Compliance with the Filtration and Disinfection
Requirements for Public Works Systems Using Surface Water Sources. Prepared by Malcolm
Pirnie, Inc. and HDR Engineering for USEPA. Contract No. 68-01-6989.
22. Wagner, R.R. 1951. "Studies on the Inactivation of Influenza Virus." Yale J. Biol. Med. pp. 288-
298.
23. Webber, W.J., Jr., and H.S. Posselt. 1972. "Disinfection." Physicochemical Processes in Water
Quality Control. W. J. Webber (editor). John Wiley & Sons, New York, NY.
24. Yahya, M.T., T.M. Straub, and C.P. Gerba. 1990a. Inactivation ofpoliovirus type 1 by Potassium
Permanganate. University of Arizona Preliminary Research Report, Tucson, AZ.
25. Yahya, M.T., Landeen, L.K., and Gerba, C.P. 1990b. Inactivation of Legionella pneumophila by
Potassium Permanganate. Environ. Technol. 11:657-662.
26. Yahya, M.T., et al. 1989. "Evaluation of Potassium Permanganate for the Inactivation of MS-2 in
Water Systems." /. Environ. Sci. Health. A34(8):979-989.
27. Young, J.S. and P.C. Singer. 1979. "Chloroform Formation in Public Water Supplies: A Case
Study." J. AWWA. 71(2):87.
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6. CHLORAMINES
The disinfectant potential of chlorine-ammonia compounds or chloramines was identified in the early
1900s. The potential use of chloramines was considered after observing that disinfection by chlorine
occurred in two distinct phases. During the initial phase, chlorine reducing compounds (i.e.,
demand) cause the rapid disappearance of free available chlorine. However, when ammonia was
present bactericidal action was observed to'continue [even though free chlorine residual was
dissipated]. The subsequent disinfection phase occurs by the action of the inorganic chloramines.
6.1 Chloramines Chemistry
Chloramines are formed by the reaction of ammonia with aqueous chlorine (i.e., HOC1). Initially,
chloramines were used for taste and odor control. However, it was soon recognized that chloramines
were more stable than free chlorine in the distribution system and consequently were found to be
effective for controlling bacterial regrowth. As a result, chloramines were used regularly during the
1930s and 1940s for disinfection. Due to an ammonia shortage during World War II, however, the
popularity of chloramination declined. Concern during the past two decades over chlorinated
organics (e.g., THM and HAA formation) in water treatment and distribution systems, increased
interest in chloramines because they form very few disinfection byproducts (DBFs).
6.1.1 Equilibrium, Kinetic, and Physiochemical Properties
Chloramines are formed from the reaction of chlorine and ammonia. The mixture that results may
contain monochloramine (NH2C1), dichloramine (NHCl2), or nitrogen trichloride (NC13). When
chlorine is dispersed in water, a rapid hydrolysis occurs according to the following reaction:
c/2 + H2o -» HOCI+H+ + cr
The equilibrium constant (Keq) at 25°C is 3.94 x 104 M"1 for this reaction. In dilute solutions at pH
greater than 3, the forward reaction is essentially complete. Hypochlorous acid (HOCI) is a weak
acid that dissociates as follows:
HOCI <^ OCr + H+ pKa = 1.6
Relative proportions of HOCI and OC1" are dependent upon pH. Both of the chlorine species in the
above reaction are powerful oxidants, capable of reacting with many substances present in water. In
aqueous solutions with pH 7.0 to 8.5, HOCI reacts rapidly with ammonia to form inorganic
chloramines in a series of competing reactions (White, 1992). The simplified stoichiometry of
chlorine-ammonia reactions are as follows:
j + HOCI -> NH2Cl + H2O (monochloramine)
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NH2Cl + HOCl
NHC12 +H2O
NHCL, + HOCl -> NC13 + H20
(dichloramine)
(nitrogen trichloride)
These competing reactions, and several others, are primarily dependent on pH and controlled to a
large extent by the chlorine:ammonia nitrogen (C^rN) ratio. Temperature and contact time also play
a role. Figure 6-1 shows the typical relationships between the chloramine species at various Cl2:N
ratios for pHs ranging from 6.5 to 8.5. This figure shows that monochloramine is predominately
formed when the applied C12:N ratio is less than 5:1 by weight. As the applied C12:N ratio increases
from 5:1 to 7.6:1, breakpoint reaction occurs, reducing the residual chlorine level to a minimum.
Breakpoint chlorination results in the formation of nitrogen gas, nitrate, and nitrogen chloride. At
C^iN ratios above 7.6:1, free chlorine and nitrogen trichloride are present. Figure 6-2 shows the
relationship between chloramine species as the pH changes (Palin, 1950). The Figure shows that
dichloramine becomes a dominant species at low pH.
2 x
DC *?
0) .P
C ^
I a
O E
HOCl + OC1-
NC1
2 4 6 8 10 12 14
Chlorine Dose, mg Cl^mg NH4-N
16
Figure 6-1. Theoretical Breakpoint Curve
To avoid breakpoint reactions, utilities should maintain a Cl2:N ratio between 3 and 5 by weight. A
ratio of 6 is actually optimum for disinfection, but it is difficult to maintain a stable operation at that
point in the breakthrough curve. Therefore, a Cl2:N ratio of 4 is typically accepted as optimal for
chloramination.
Furthermore, over a period of a day or so, without any modification of pH or Cl2:N ratio,
monochloramine will degrade slowly to dichloramine to a ratio of 43 percent NHaCl to 57 percent
NHCh. Dichloramine is relatively unstable in the presence of HOCl; therefore, pure solutions of this
form of monochloramine are difficult to generate and maintain.
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6. CHLORAMINES
U
•a
CD
!S
I
U
Source: Palin, 1950.
Figure 6-2. Distribution Diagram for Chloramine Species with pH
6.2 Generation
Chloramines are formed by the reaction of hypochlorous acid and ammonia according to the
equations described in Section 6.1. Table 6-1 summarizes the theoretical doses of chlorine and
ammonia based on these formulas. Monochloramine is the preferred chloramine species for use in
disinfecting drinking water because of taste and odor problems associated with dichloramine and
nitrogen trichloride. To ensure that these compounds are not formed, common practice was to limit
the chlorine to ammonia ratio to 3:1. However, because of problems such as nitrification and biofilm
growth, which can be caused by excess ammonia, current practice is to use a C12:N ratio in the range
of 3:1 to 5:1, with a typical value of 4:1.
Table 6-1. Chlorine Dose Required for NH3 - CI2 Reaction
Reaction
Monochloramine (NHsCI)
Dichloramine (NHCb)
Nitrogen Trichloride (NCb)
Nitrogen (N2)
Nitrate (NOs)
Free residual reaction
mg Cl2/mg JNJJ3
4.2
8.4
12.5
6.3
16.7
9
Source: AWWA and ASCE, 1990.
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6. CHLORAMINES
The rate of reaction of monochloramine formation is sensitive to pH. Table 6-2 shows the calculated
reaction times for monochloromine formation at 25°C, and at a chlorine:amonia ratio of 3:1 (White,
1992).
Table 6-2. Time to 99 Percent Conversion of Chlorine to Monochloramine
PH
Time (seconds)
2
4
7
8.3
12
421
147
0.2
0.069
33.2
6.2.1 Chlorine Feed Facilities
Table 6-3 summarizes commonly used methods of chlorine addition, including their safety
precautions and costs.
Table 6-3. Methods of Chlorine Addition
Method
Description
Safety precautions
Costs
Gaseous Gas delivered in containers
chlorine ranging in size from 150 Ib
cylinders to 90 ton rail cars. One
ton cylinders are commonly used.
Feed equipment consists of
solution water pump/ejector to
create vacuum and automatic
orifice control to meter the gas.
Gas can be drawn directly from
storage container or be generated
by an evaporator from liquid
withdrawn from the container. A
schematic of gaseous chlorine
feed system is shown in
Figure 6-3.
Sodium Sodium hypochlorite can be
hypochlorite purchased bulk in quantities
ranging from 55 gal drums to
4,500 gal truck loads. Bulk loads
can be stored in fiberglass or
plastic tanks. Solution is fed
directly into the process stream. A
schematic of typical hypochlorite
feed system is shown in Figure 6-
4.
Gaseous chlorine is classified The cost per
by the Uniform Fire Code as pound of
an oxidizing, highly toxic, liquid chlorine
compressed gas. New is in the range
gaseous chlorine facilities of $0.08 to
should be designed with $0.20 per
enclosures and air scrubbers pound
to capture and neutralize any depending on
gas that leaks. Risk the quantity
management prevention plans purchased.
should be prepared. Personnel
safety equipment and training
should be provided for
operators.
Hypochlorite solution is toxic
and classified as hazardous.
Storage facilities should be
designed with secondary
containment.
Typical
chemical cost
is $0.60 to
$1.00 per
pound CI2.
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6.2.2 Ammonia Feed Facilities
Ammonia feed facilities can be located on-site at the water treatment plant or at remote locations in
the distribution system (Dennis et al., 1991). Most ammonia feed facilities use either gaseous
(anhydrous ammonia) or liquid (aqueous) ammonia. Though anhydrous ammonia is a gas at ambient
temperature and pressure, it is commonly stored and transported as a liquid in pressure vessels. In
this phase, ammonia is highly soluble in water. Storage facilities and handling equipment should be
kept dry (Dennis et al., 1991).
Horizontal Manifold
Assembly
Vacuum Regulator
Gas Feeder
Flexible Connector
Chlorine Cylinder
(150 Ib, Ton, Tanker)
Controller
Vacuum Pipe
Eductor
Water Supply
Chlorine Solution
Figure 6-3. Gaseous Chlorine Feed System
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Chemical
Storage Tank
Calibration
Column
Tank
Shut-off
Valve
Relief
Valve
Pulsation
Darapner
Back
Pressure
Valve
Shut-
off
Valve
-CX}-
Strainer
Metering
Pump
Flow to
Discharge
Point
Figure 6-4. Hypochlorite Feed System
5.2.2.1 Anhydrous Ammonia
Anhydrous ammonia is stored in portable cylinders or stationary tanks. Portable cylinders are similar
to chlorine cylinders and are available in 100, 150, and 800 Ib sizes (Dennis et al., 1991). The
cylinders are rated for a minimum service pressure of 480 psi. Stationary tanks are typically 1,000
gallon vessels that can be used on-site. These tanks are refilled by tanker trailers. The storage tanks
can be located indoors or outdoors. Since each tank has a minimum working pressure of 250 psi
(valves and fittings on the tanks are rated for 300 psi), a tank stored outdoors should have protection
from extreme temperatures (greater than 125°F and less than 28°F) (Dennis et al., 1991). In warmer
climates, an outdoor tank should be painted white and protected from sunlight. In colder climates,
the tank should be wrapped with heat tape to prevent impairment of the ammonia vaporization.
Anhydrous ammonia is applied using an ammoniator. An ammoniator is a self-contained modular
unit with a pressure reducing valve, gas flow meter, feed rate control valve, and miscellaneous piping
for controlling the flow of ammonia. Automatic paced ammoniators are available. An evaporator is
used when large quantities of ammonia gas are needed. An anti-siphon valve or check valve should
be used to prevent water from entering the ammoniator.
Anhydrous ammonia is usually applied by direct feed or solution feed. The direct feed method is
typically used when the process stream has a low pressure and the ammonia feed rate is less than
1,000 Ib per day (i.e., maximum rated feed capacity). Ammonia is drawn from the storage tank
under high pressure (e,g., 200 psi), and injected directly into the process stream at a low pressure of
15 psi. The tank pressure is first reduced by a pressure reducing valve to approximately 40 psi, and
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6. CHLORAMINES
then by another pressure reducing valve in the ammoniator. Typical application points are at open
channels and basin facilities. Figure 6-5 is a schematic of a direct anhydrous ammonia feed system.
The solution feed method is typically used where direct feed systems are not adequate (e.g., ammonia
feed rate is greater than 1000 Ib/day or where the process stream pressure is high) (Dennis et al.,
1991). This type of application is similar to the chlorine vacuum feed system. The supply tank
pressure is reduced by a pressure reducing valve to create a vacuum. An eductor is used to withdraw
ammonia from the ammoniator where the ammonia is dissolved into a side water stream and pumped
into the process stream. Solution feed ammoniators are available up to 4,000 Ib/day capacities and
can operate at discharge pressures up to 100 psi (Dennis et al., 1991). Softened water (i.e., hardness
less than 29 mg/L as CaCO3) is required for the carrier stream. Otherwise, the ammonia addition will
precipitate scale that may plug the eductor and application point. Figure 6-6 shows a schematic of a
solution feed system.
6.2.2.2 Aqueous Ammonia
Aqueous ammonia is produced by dissolving anhydrous ammonia into deionized or softened water.
This form of ammonia is shipped in cargo trucks or polyethylene lined steel drums. Plastic drums
are not recommended since they tend to lose their shape under the slight pressure exerted by the
aqueous ammonia. Aqueous ammonia is stored in low pressure tanks, typically steel or fiberglass.
Since excessive temperatures will cause ammonia gas to vaporize, each storage tank should be
equipped with a water trap or ammonia scrubber to keep vapors from escaping to the atmosphere.
Flow Signal
I Anhydrous NH3 tank trailer I
• 77A
in parallel J
Custom gas diffuser
(maximum discharge pressure
= 15 psi for direct-feed
ammoniators)
Ammoniation building
Source: Montgomery, 1985.
Figure 6-5. Anhydrous Ammonia Direct Feed System
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6. CHLORAMINES
Flow Signal
^*—~ Control
Control System
Anhydrous NH3 tank trailer)
Evaporator V-Notch
x (if necessary) ammoniator
Ammoniatlon Building
Utility Water
1 r
Brine waste
during regeneration
Pump
Solution dil't'uscr
(- 150 psi maximum
discharge pressure)
Source; Montgomery, 1985.
Figure 6-6. Anhydrous Ammonia Solution Feed System
Aqueous ammonia feed systems are similar to other liquid chemical feed systems. They require a
storage tank, chemical metering pump, relief valve, pulsation dampener, flow meter, and
backpressure valve. Typically, the feed pumps are positive displacement or progressive cavity type
metering pumps. The feed pumps should be placed fairly close to the storage tank to minimize
chances of ammonia vaporization in the piping (Dennis et al., 1991). The pump should be designed
to compensate for changes in ambient temperatures, different aqueous ammonia solutions, and
changes in the chlorine-to-ammonia ratio (Skadsen, 1993). When aqueous ammonia is applied to
water, complete mixing is required for the ammonia to react with chlorine in the water to reduce the
formation of dichloramine and nitrogen trichlorine. Figure 6-7 shows a schematic of an aqua
ammonia feed system.
Metropolitan Water District of Southern California (MWDSC) uses aqueous ammonia at its
chloramination facility. Ammonia is stored in unlined tanks and pumped to the ammoniator with
progressive cavity pumps. During startup of its aqueous ammonia feed system, MWDSC
experienced complete pump failures. Based on MWDSC's experience, EPDM rotors with adequate
quality chromed finish stators are recommended for progressive cavity pumps. A mechanical seal is
also recommended instead of a packing box to reduce the possibility of ammonia leaks (Skadsen,
1993). MWDSC also later installed special blow-offs and strainers in the feed pump suction line to
reduce plugging at the magnetic flow meters. The pump problems prior to startup led MWDSC to
install an alternative, redundant ammonia feed system. A pressurized system was designed to feed
aqueous ammonia by pressurizing the ammonia tanks and by-passing the pump.
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6. CHLORAMINES
Control
System
Aqua NH3
on-site
storage tank
Flow Signal
Flow
Meter
| | Back-
pressure
valve
Relief Valve
•B-
Pulsation
t 'Dampener
Metering
Pump
Solution diffuser
(- ISOpsi maximum
discharge pressure)
Source: Montgomery, 1985.
Figure 6-7. Aqua Ammonia Feed System
A 5.5 gpm flexible impeller centrifugal pump with a recirculation loop back to the storage tank
regulates the back pressure on the by-passed feed pump. This alternative, redundant method proved
to be reliable and economical. In addition, it provided a stable feed rate and required little
maintenance (Skadsen, 1993).
6.2.2.3 Piping and Valving
For anhydrous ammonia, the typical piping materials for both direct and solution feed systems are
stainless steel, PVC, and black iron (Dennis et al., 1991). Stainless steel or black iron pipe is used in
the high pressure (i.e., greater than 15 psi) portions of the feed system. PVC pipe is used only in the
low pressure portion of the feed system, after the ammoniators.
For aqueous ammonia, PVC piping should be used due to. the corrosive nature of aqueous ammonia
(Dennis et al., 1991).
6.2.2.4 Safety Provisions for Chloramine Generation Facilities
A chloramination facility should include some safety provisions to prevent the formation of nitrogen
trichloride and the vaporization of ammonia at ambient temperatures. The possible formation of
nitrogen trichloride at a chloramination facility should be considered when selecting sites for the
ammonia and chlorine storage facilities.
Dennis et al. (1991) provides detailed information about safety provisions for chlorarnine facilities.
Chlorine gas and ammonia gas should never be stored in the same room. The ammonia gas
application points should be located at least 5 feet away from chlorine feed solution lines. Anhydrous
ammonia is lighter than air, so any leaking vapor will rise quickly. Under pressure, anhydrous
ammonia is a liquid. Great amounts of heat are absorbed when the pressurized liquid reverts to a gas.
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If the storage tanks and/or chemical feed equipment are installed indoors, ventilation and vapor
detection devices should be located at high points in the room. The ventilation rates will vary
depending on the appropriate regulatory agency's requirements. Typically, a minimum of 6 room
volume changes per minute is recommended.
Ammonia gas storage tanks should be protected from direct sunlight or direct sources of heat (i.e.,
greater than 125°F) to avoid pressure increases in the tank (Dennis et al., 1991). Otherwise,
ammonia gas may be released into the atmosphere through the pressure relief valves. In warm
regions, outdoor tanks should be covered with a shelter or outfitted with a temperature control
sprinkler system. Where fugitive emissions of ammonia are a concern, fume control may be
required. If the accidental release from a storage container is a concern, an emergency scrubber
system similar to a chlorine gas scrubber system should be considered.
6.3 Primary Uses and Points of Application
Monochloramine is used in drinking water treatment for disinfection and nuisance organism control.
Points of application are based on treatment objectives and contact time disinfection requirements.
6.3.1 Primary Uses
6.3.1.1 Disinfection
The primary use of monochloramine in water systems is as a secondary disinfectant for maintaining a
residual in the distribution system. Chloramines are a good choice for secondary disinfectant because
of the following potential benefits:
* Chloramines are not as reactive with organics as free chlorine in forming THMs.
• The monochloramine residual is more stable and longer lasting than free chlorine or chlorine
dioxide, providing better protection against bacterial regrowth in systems with large storage tanks
and dead-end water mains.
• The monochloramine residual has been shown to be more effective in controlling biofilms
because of its superior ability to penetrate the biofilm. Controlling biofilms also tends to reduce
coliform concentrations and biofilm induced corrosion.
• Because Chloramines do not tend to react with organic compounds, many systems will experience
less incidence of taste and odor complaints when using chloramines.
Water systems in Indiana and Virginia found that conversion from free chlorine to monochloramine
as the secondary disinfectant significantly reduced coliform concentrations in the distribution system
(Norton and LeChevallier, 1997).
The normal dosage range for monochloramine is in the range of 1.0 to 4.0 mg/L. The minimum
residual of monochloramine in the distribution system is typically regulated at 0.5 mg/L (Texas
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6. CHLORAMINES
Natural Resource Conservation Commission). For prevention of nitrification in a distribution system,
a minimum monochloramine dosage of 2.0 mg/L is recommended (Norton and LeChevallier, 1997).
6.3.1.2 Nuisance Organism Control
Cameron et al. (1989a) investigated the effectiveness of monochloramine to control the Asiatic clam
in both the juvenile and adult phases. The adult Asiatic clam was found to be much more resistant to
monochloramine than the juvenile form. Monochloramine was found to be the most effective for
controlling the juvenile Asiatic clam in terms of LT50 (time required for 50 percent mortality).
Monochloramine doses used to control the juvenile Asiatic clam range from 1.2 to 4.7 mg/L. Further
research showed that the effectiveness of monochloramine increased greatly as the temperature
increased (Cameron et al., 1989b).
6.3.2 Points of Application
The formation of monochloramine can be accomplished by first adding ammonia and then chlorine,
or vice versa. Ammonia is added first where formation of objectionable taste and odor compounds
caused by the reaction of chlorine and organic matter are a concern. However, most drinking water
systems add chlorine first in the treatment plant in order to achieve the required concentration and
contact time (CT) to meet EPA's SWTR disinfection requirements. Typically, the point of ammonia
addition is selected to "quench" the free chlorine residual after a target period of time based on
optimizing disinfection versus minimizing DBF formation.
Because the germicidal effectiveness of monochloramine is a factor of 200 less than for free chlorine,
extremely long contact times are required for monochloramine to meet EPA disinfection CT
requirements. Therefore, if ammonia is added first, a means of ensuring that CT requirements are met
must be developed.
6.3.2.1 Impact on Other Treatment Processes
Monochloramine addition impacts other processes at the water treatment facility. These impacts
include:
• Ammonia used in the chloramination process can provide nutrient ammonia for nitrifying
bacteria growth in the distribution system, which can cause increased nitrate levels in the
finished water where systems do not normally test for nitrate.
• Imbalances in chlorine and ammonia concentrations (in greater than an 8 to 1 ratio) can cause
breakpoint chlorination reactions to occur when encountered in distribution system
• Monochloramine addition upstream of filters will reduce biological growth on filters. This
has a favorable impact on the filters by keeping them clean and reducing the backwash
frequency. It also has the undesirable impact of reducing BDOC removal in the filters when
the filters are run in a biological mode.
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6. CHLORAMINES
The reader is referred to EPA's Microbial and Disinfection Byproduct Simultaneous Compliance
Guidance Document (expected to be available in 1999) for additional information regarding the
interaction between oxidants and other treatment processes.
6.4 Pathogen Inactivation and Disinfection Efficacy
Chlexamination of drinking water has been practiced in the United States for nearly 80 years. In
addition to achieving disinfection, chloramines have been used by the Denver Water Department for
controlling tastes and odors since 1918 (Hazen and Sawyer, 1992). Chloramination has also been
found to provide a more stable residual in water distribution system. However, because of its
relatively weak disinfecting properties for inactivation of viruses and protozoa pathogens, it is rarely
used as a primary disinfectant, and then only with long contact times.
6.4.1 Inactivation Mechanisms
The mechanisms by which chloramines inactivate microorganisms have been studied to a lesser
degree than chlorine. A study of inactivation of E. coli by chloramines concluded that
monochloramine readily reacts with four amino acids; cysteine, cystine, methionine and
tryptophan (Jacangelo et. al, 1987). The mechanism of inactivation for chloramine is therefore
thought to involve inhibition of proteins or protein mediated processes such as respiration.
Jacangelo further concluded that because of the inconsistency in rate of inactivation
monochloramine should have "multiple hits" upon bacterial cells before cell death.
Few studies have been performed to determine the mechanism for viral inactivation. The initial site
for destruction of bacteriophage f2 involved the RNA fragment (Olivieri et al., 1980). However, the
primary mechanism for poliovirus inactivation by chloramines involved the protein coat (Fujioka et
al., 1983). Similar to free chlorine, the mechanism of viral inactivation by chloramine may be
dependent on factors such as virus type and disinfectant concentration.
6.4.2 Environmental Effects
Several studies have been performed to determine the effect of pH, temperature, and organic and
inorganic compounds on the disinfection effectiveness of chloramines. Following is a summary of
the affect these parameters have on pathogen inactivation.
6.4.2.1 pH
The effect of pH on disinfection has more to do with the organism than with the disinfectant;
however, pH also impacts disinfection efficiency by controlling the chloramine species distribution.
Studies have indicated that the disinfection efficacy of monochloramine and dichloramine are not
equal. One study showed that the bactericidal properties of dichloramine were superior to that of
monochloramine (Esposito, 1974). However, pH may be a compounding factor because changes in
pH may alter the physiological response of the organism (Hoff and Geldreich, 1981). Other studies
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6. CHLORAMfNES
have shown that monochloramine is superior to dichloramine with regard to virucidal ability (Dorn,
1974; Esposito, 1974; Olivieri et al., 1980). Some evidence suggests that solutions containing
approximately equal concentrations of monochloramine and dichloramine may be more
microbiocidal than those containing only monochloramine or dichloramine (Weber and Levine,
1944).
6.4.2.2 Temperature
Similar to most of the disinfectants discussed in this report, the bactericidal and viral inactivation
efficiency of chloramine increases with increasing temperature. Moreover, the efficiency
dramatically decreases under conditions of high pH and low temperature. For example, the
inactivation of E. coli. is approximately 60 times slower at pH 9.5 and temperatures of 2 and 6°C
than at pH 7 and temperatures between 20 and 25°C (Wolfe et al., 1984). Similar results were
obtained for poliovirus 1 inactivation (Kelley and Sanderson, 1958).
6.4.2.3 Organic Nitrogen and Other Compounds
In addition to ammonia, free chlorine reacts with organic nitrogen compounds to form a variety of
organic chloramines. These organic chloramines are undesirable byproducts because they exhibit
little or no microbiocidal activity (Feng, 1966). Studies have indicated that chlorine binds to amine-
containing compounds more rapidly than to ammonia (Weil and Morris, 1949; Morris, 1967;
Margerum et al., 1978) and that chlorine can be transferred from inorganic chloramines to amine-
containing compounds (Margerum et al., 1978; Isaac and Morris, 1980)
Several other reactions may occur which divert chlorine from the formation of chloramines. These
reactions can include oxidation of iron, manganese, and other inorganics such as hydrogen sulfide
(Hazen and Sawyer, 1992).
6.4.3 Disinfection Efficacy
Chloramines are relatively weak disinfectants for virus arid protozoa inactivation. As a consequence,
it is extremely difficult to meet the SWTR CT criteria for primary disinfection of Giardia and viruses
using chloramines because very long detention times are needed. However, given the ability of
chloramines to provide a stable residual, this form of disinfection appears to be feasible for
secondary disinfection protection against microbial growth in distribution systems. The following
paragraphs describe the disinfection efficiency of chloramines in terms of bacteria, virus, and
protozoa inactivation.
6.4.3.1 Bacteria Inactivation
A series of comprehensive experiments was initiated in the mid 1940s to determine the relative
bactericidal effectiveness of free chlorine and inorganic chloramines. Results from these
experiments showed conclusively that under relatively demand-free, laboratory-controlled
conditions, free chlorine inactivated enteric bacteria much faster that chloramines (Wattie and
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Butterfield, 1944). In this experiment, a monochloramine concentration of 0.3 mg/L required 240
minutes of contact time for 3-log inactivation of E. coli whereas exposure to 0.14 mg/L free C12
required only 5 minutes to achieve the same level of inactivation at the same temperature and pH.
6.4.3.2 Virus Inactivation
According to reports written by Kabler et al. (1960) and the National Research Council (1980), all
studies conducted prior to 1944 that compared virucidal potency of free and combined chlorine were
inaccurate because the experiments failed to clearly differentiate between free and combined forms
of chlorine and because high-chlorine-demand water was used in the experiments.
The majority of the experiments conducted after the mid-1940s has shown that inorganic chloramines
require much higher concentrations and considerably longer contact times than free chlorine to
achieve comparable levels of virus inactivation. Experiments showed that contact times between 2
and 8 hours were required at concentrations between 0.67 to 1.0 mg chloramines to achieve greater
than 2-log inactivation of poliovirus 1 (Mahoney and MK500), poliovirus 2 (MEF), poliovirus 3
(Sackett), coxsackievirus Bl, and coxsackievirus B5 (EA 80) (Kelley and Sanderson, 1958 and
1960). In contrast, 0.2 to 0.35 mg/L free Clj required 4 to 16 minutes of contact time to achieve
comparable levels of inactivation under the same conditions.
6.4.3.3 Protozoa Inactivation
Of the three predominant forms of pathogens (i.e., bacteria, viruses, and protozoan [oo]cysts), studies
have shown that protozoan [oo]cysts are usually the most resistant to all forms of disinfection.
Studies have indicated that free chlorine is a more effective disinfectant than chloramines for
[oo]cyst inactivation (Chang and Fair, 1941; Chang, 1944; Stringer and Kruse, 1970). Chloramine
concentrations of 8 mg/L were required for 2-log inactivation of Entamoeba histolytica cysts whereas
only 3 mg/L of free chlorine was required to obtain the same degree of inactivation (Stringer and
Kruse, 1970). Contact times for both disinfectants were 10 minutes.
6.4.3.4 CT Values
CT values for achieving Giardia cyst and virus inactivation using chloramines are shown in
Table 6-4 and Table 6-5, respectively. Values contained in these tables were obtained from the
Guidance Manual for Compliance with Filtration and Disinfection Requirements for Public Water
Systems Using Surface Water Sources (AWWA, 1991).
CT values shown in Table 6-4 are based on disinfection studies using in vitro excystation of Giardia
lamblia. CT values shown in Table 6-5 were based on data using preformed chloramines at pH 8. No
safety factor was applied to laboratory data used to derive the CT values shown in Table 6-4 and
Table 6-5 since chloramination conducted in the field is more effective than using preformed
chloramines, since monochloramine tends to degrade with time and some free chlorine is present
when forming chloramines which enhances the inactivation process.
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6. CHLORAMINES
Table 6-4. CT Values for Giardia Cyst Inactivation Using Chloramines
Temperature (°C) (mg»min/L)
Inactivation
0.5-log
1-log
1 .5-log
2-log
2.5-log
3-log
5
365
735
1,100
1,470
1,830
2,200
10
310
615
930
1,230
1,540
1,850
15
250
500
750
1,000
1,250
1,500
20
185
370
550
735
915
1,100
25
125
250
375
500
625
750
Source: AWWA, 1991.
Values shown in this table are based on a pH range between 6 and 9.
Table 6-5. CT Values for Virus Inactivation Using Chloramines
Temperature (°C) (mg1
Inactivation
2-log
3-log
4-log
5
857
1,423
1,988
10
643
1,067
1,491
15
428
712
994
•min/L)
20
321
534
746
25
214
356
497
Source: AWWA, 1991.
6.5 DBP Formation
The effectiveness of chloramines to control DBP production depends upon a variety of factors,
notably the chlorine-to-ammonia ratio, the point of addition of ammonia relative to that of chlorine,
the extent of mixing, and pH.
Monochloramine (NH2C1) does not produce DBFs to any significant degree, although some
dichloroacetic acid can be formed from monochloramine and cyanogen chloride formation is greater
than with free chlorine (Jacangelo et al., 1989; Smith et al., 1993; Cowman and Singer, 1994). The
inability to mix chlorine and ammonia instantaneously allows the free chlorine to react before the
complete formation of chloramines. In addition, monochloramine slowly hydrolyzes to free chlorine
in aqueous solution. Therefore, halogenation reactions occur even when monochloramine is formed
prior to addition in the treatment process (Rice and Gomez-Taylor, 1986). The closer the
chlorine:ammonia ratio is to the breakpoint, the greater the formation of DBFs (Speed et al., 1987).
In addition to controlling the formation of DBFs, chloramination results in lower concentrations of a
number of the other specific organic halides generated from free chlorine, except for cyanogen
chloride (Krasner et al., 1989; Jacangelo et al., 1989). Increased production of cyanogen chloride is
observed when monochloramine is used as a secondary disinfectant instead of free chlorine.
The application of chloramines results in the formation of chlorinated organic material, although it
occurs to a much lesser degree than from an equivalent dose of free chlorine. Little is known about
the nature of these byproducts, except that they are more hydrophilic and larger in molecular size
than the organic halides produced from free chlorine (Jensen et al., 1985; Singer 1993).
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6. CHLORAMINES
6.6 Status of Analytical Methods
6.6.1 Monitoring of Chloramines
There has been little development of analytical procedures for selective determination of
monochloramines (Gordon, et al., 1992). Typically, the methods used for chlorine residual
measurement are adapted for chloramine measurement. The DBPR promulgated on December 16,
1998 (63 FR 69390) establishes three analytical methods that are acceptable for measuring residual
chlpramines (combined chlorine). These methods are presented in 40 CFR § 141.131 (c) and include:
• Amperometric Titration (Standard Method 4500-C1D and ASTM Method D 1253-86);
• DPD Ferrous Titrimetric (Standard Method 4500-C1F); and
• DPD Colorimetric (Standard Method 4500-C1 G).
If approved by the State, systems may also measure chloramines by using DPD colorimetric test
kits.
6.6.1.1 Amperometric Tltrations
The amperometric titration method is utilized extensively in water treatment laboratories (Gordon, et
al., 1992). This method capable of differentiating the three most common forms of chlorine, namely
chlorine/hypochlorous acid/hypochlorite ion, monochloramine, and dichloramine, as long as the
combined forms are not present in concentrations greater than about 2 mg/L (as Ck). For higher
concentrations, dilution of the samples is required, but differentiation is still possible (Aoki, 1989).
The amperometric titration method is a standard of comparison for the determination of free or
combined chlorine. This method is not greatly affected by common oxidizing agents, temperature
variations, turbidity, and color (Standard Methods, 1995). Amperometric titration requires a greater
degree of skill than colorimetric methods. The differentiation of free chlorine, monochloramine, and
dichloramine is possible by control of potassium iodide (KI) concentration and pH during the
analysis.
Several methods are commonly used to measure chlorine species utilizing the amperometric titration
including forward and back titration (Gordon, et al., 1992). The lower limit of detection of these
methods varies depending on the instrumentation used and type of water sample analyzed. The lower
limit of detection for commercial amperometric titrating equipment is about 30 )j,g/L as Ch (Sugam,
1983).
Table 6-6 shows the working range, expected accuracy and precision, operator skill level required,
interferences, and current status for amperometric method monochloramine analysis comparison.
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6. CHLORAMfNES
6.6.1.2 Colorimetric Methods
Over the years, numerous colorimetric methods have been developed to measure free and combined
chlorine in aqueous solutions (Gordon, et al., 1992). Not many of these methods would be
recommended as the method of choice. Two of the colorimetric methods listed in Standard Methods
(1995), are DPD methods. In addition, the colorimetric LCV method modified by Whittle and
Lapteff (1974) can be used to measure free and combined chlorine species.
The DPD methods (ferrous titration and colorimetric) are operationally simpler for determining free
chlorine than the amperometric titration (Standard Methods, 1995). Procedures are given for
estimating separate monochloramine and dichloramine fractions, as well as combined chlorine
fractions.
The LCV method modified by Whittle and Lapteff modifies the discontinued Standard Method for
LCV. The maximum chlorine concentration that can be determined by this modified procedure,
without dilution of the sample, is 10 mg/L as C12 (Whittle and Lapteff, 1974).
See Table 6-6 for the working range, expected accuracy and precision, operator skill level required,
interferences and current status for colorimetric method monochloramine analysis comparison.
6.6.2 Disinfectant Interferences
Interferences to free chlorine may impact the measurement of monochloramine since the methods use
the free chlorine level in the determination of monochloramine. Many strong oxidizing agents
interfere in the measurement of free chlorine in all monochloramine methods, including bromine,
chlorine dioxide, iodine, permanganate, hydrogen peroxide, and ozone. However, the reduced form
of these compounds (i.e. bromide ion, chloride ion, iodide ion, manganous ion, and oxygen) do not
interfere. Reducing agents such as ferrous compounds, hydrogen sulfide, and oxidizable organic
matter generally do not interfere (Standard Methods, 1995).
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onochloramine" Analytical Methods
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6. CHLORAMIWES
6.6.2.1 Amperometric Titrations
The amperometric titration methods are unaffected by dichloramine concentrations in the range of 0
to 9 mg/L as Cla in the determination of free chlorine. Nitrogen trichloride, if present, may react
partially as free chlorine. The amperometric method will measure organic chloramines as free
chlorine, monochloramine, or dichloramine, depending upon the activity of chlorine in the organic
sample (Standard Methods, 1995). Dichloramine may also interfere with the measurement of both
monochloramine and free chlorine (Marks, et al., 1951). The presence of iodide ion can be a severe
problem if the titrator glassware is not washed carefully between determinations (Johnson, 1978).
Manganese dioxide, a common interference in most chlorine analytical procedures, does not interfere
in the amperometric measurement of free chlorine (Bongers et al., 1977). However, because of its
reaction with iodide ion, added during analysis, manganese dioxide does interfere with the
amperometric measurement of combined forms of chlorine such as monochloramine (Johnson,
1978).
6.6.2.2 Colorimetric Methods
Sample color and turbidity may interfere in all colorimetric procedures. In the DPD colorimetric
methods, high concentrations of monochloramine interfere with free chlorine determination unless
arsenite or thioacetimide are added. In addition, the DPD methods are subject to interference by
oxidized forms of manganese unless compensated for by a blank (Standard Methods, 1995). The
DPD methods are unaffected by dichloramine concentrations in the range of 0 to 9 mg/L as C\2 in the
determination of free chlorine. Nitrogen trichloride, if present, may react partially as free chlorine.
The extent of this interference in the DPD methods does not appear to be significant (Standard
Methods, 1995).
In the LCV colorimetric method, Whittle and Lapteff (1974) reported that dichloramine did not
interfere with the monochloramine measurement.
6.6.3 Chloramine Monitoring for Systems Using Chloramines
Pursuant to 40 CFR §141.132(c)(l), community water systems and non-transient non-community
water systems that use chloramines, must measure the residual disinfectant level at the same points in
the distribution system, at the same time, and at the same frequency (based on population served) as
total coliforms are sampled, as specified in 40 CFR §141.21. These systems may use the results of
residual disinfectant concentration sampling conducted under §141.74(b)(6)(i) for unfiltered systems
or §141.74(c)(3)(i) for systems which filter, in lieu of taking separate samples. No reduced
monitoring allowances exist for these systems.
Compliance with the MRDL of 4.0 mg/L (as chlorine) is based on a running annual arithmetic
average, computed quarterly, of monthly averages of all samples collected by the system under
§141.132(c)(l). If the average quarterly averages covering any consecutive four-quarter period
exceeds the MRDL, the system is in violation of the MRDL and must notify both the public, pursuant
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6. CHLORAMINES
to §141.32, and the State, pursuant to §141.134. Where systems switch between the use of chlorine
and chlpramines for residual disinfection during the year, compliance is determined by including
together all monitoring results of both chlorine and chloramines in calculating compliance.
6.7 Operational Considerations
•i ,j <
The purpose of this section is to address operational considerations in the use of chloramines in
drinking water treatment. Specifically, the following topics are addressed below: the conversion of
chloramination from chjorination; the potential operational impacts from chloramination disinfection;
and special considerations for chloramination facilities. For a more detailed discussion of
chloramine disinfection, refer to "Optimizing Chloramine Treatment" by Kirmeyer et, al. 1993.
6.7.1 Conversion to Chloramination from Chlorination
6.7.1.1 Planning
Project planning and preparation are essential to ensure an efficient changeover, maintain a
dependable and safe system, and preserve the public confidence in the water purveyor (Skadsen,
1993). Planning and preparation should consider the following aspects:
• Raw water composition and suitability to chloramination;
• Treatment plant and distribution system attributes and monitoring program;
• Employee training;
• Public notification and education; and
• Environmental affects from chloraminated water.
6.7.1.2 Preliminary Analysis
A bench scale study is necessary to identify the water characteristics and to determine if
chloramination is suitable. White (1992) describes some of the study objectives and variables to
consider. The reaction time to form free chloramine residuals varies for each water source since the
reaction rate between chlorine and ammonia nitrogen depends on the water's temperature and pH of
the water. The reaction rate is also affected by the chlorine and ammonia nitrogen concentrations.
To properly control the reaction time between chlorine and ammonia, the study should use different
chlorine:ammonia nitrogen ratios, ammonia feed doses, and contact times.
,ji
The amount of ammonia required for chloramine residual disinfection depends on the following
factors (Dennis et al., 1991):
• Organic nitrogen in the water;
• Ammonia residual desired in the distribution system; and
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6. CHLORAMINES
• Chloramine residual type and concentration required in the distribution system.
If there is organic nitrogen in the untreated water, the amount of supplemental ammonia required
should be carefully determined by subtracting the background ammonia from the desired dose. The
dose should also consider the amount of ammonia residual desired in the distribution system. For
residual disinfection, approximately 1 to 2 mg/L of ammonia is required (Dennis et al., 1991).
For each specific water, a breakpoint curve should be developed to determine the chloramine residual
type required. Monochloramine residuals are preferred for most water distribution systems.
Dichloramine and nitrogen trichloride residuals may cause taste and odor problems when
concentrations exceed 0.8 mg/L or 0.02 mg/L, respectively. Monochloramines are primarily formed
when the theoretical chlorine to ammonia dose ratio is less than 5 to 1 (by weight ratio) and the pH is
greater than 7.0 (Dennis et al., 1991). The chloramine residual concentration leaving the treatment
plant will vary depending on the size of the distribution system and the chloramine demand exerted
by the system. Typical chloramine residuals range from 1 to 4 mg/L (Dennis et al., 1991).
6.7.1.3 The Metropolitan Water District (MWDSC)
MWDSC of Southern California converted from free chlorine to chloramine disinfection in 1985 to
assist its 27 member agencies in complying with the EPA's total trihalomethane regulation.
MWDSC serves approximately 15 million people and operates five treatment plants, with a
combined capacity of 1,670 MOD. Raw water is taken from two sources: the Colorado River and
California state project water.
Prior to the changeover, MWDSC performed extensive investigations into the chemical,
microbiological, and engineering aspects of chloramine disinfection. To prepare for the changeover,
MWDSC coordinated the efforts among its treatment plants, distribution system reservoirs,
laboratory personnel, and management. A formal request for approval to use chloramines as a
disinfectant was submitted to the California State Department of Health Services. Next, a series of
workshops was held on the engineering, chemical, and microbiological aspects of chloramine
disinfection. MWDSC also prepared a manual for the type of chloramination application method and
ammonia form selected. Information in the manual included the feed equipment information, project
specifications, piping layouts, preliminary analysis, and safety and maintenance issues.
It was essential to notify specific sectors of the public that could be affected by the use of
chloramines. MWDSC made its customers aware of the changeover and kept them apprised of the
options for preventing adverse reactions through an extensive notification program that involved
state and county health departments, appropriate interest groups, and the media.
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6. CHLORAMINES
6.7.2 Potential Operational Impacts from Chloramination
Disinfection
6.7.2.1 Pretreatment
Ammonia in excess of the required chlorine can promote the growth of nitrifying bacteria in filter
beds (i.e., rapid sand filters) (White, 1992). The excess ammonia acts as a nutrient and causes the
growth of nitrifying bacteria, which convert the excess ammonia to nitrates and nitrites. Excessive
levels of nitrate in drinking water have caused serious illness and sometimes death in infants under
six months of age. The symptoms include shortness of breath and blueness of skin [40 CFR
§141.32(e)(20)]. Prior to designing a chloramination facility, the amount of ammonia naturally
occurring in the raw water should be determined. The required ammonia dosage would then be
based on the anticipated naturally occurring ammonia levels.
A chloramine residual concentration should also be maintained in the discharge stream from the
filters. An adequate residual concentration would be between 0.5 to 1 mg/L chloramine (White,
1992).
6.7.2.2 Nitrification
Nitrification in chloraminated drinking waters is usually partial. Partial nitrification occurs when the
chloraminated water has excess ammonia present in the distribution system (Skadsen, 1993). Partial
nitrification can have various adverse effects on water quality, including a loss of total chlorine and
ammonia residuals and an increase in heterotrophic plate count (HPC) bacteria concentration. The
excess ammonia encourages the growth of nitrifying bacteria that convert ammonia to nitrates. An
intermediate step in this conversion results in a small amount of nitrite being formed. Research has
shown that a chlorine demand of 5 mg/L is exerted by 1 mg/L of nitrite (Cowman and Singer, 1994).
The nitrites rapidly reduce free chlorine, accelerate decomposition of chloramines, and can interfere
with the measurement of free chlorine (Skadsen, 1993). Valentine (1998) found that the decay of
monochloramine was increased (from a second order rate constant of 0.07 to 0.106) by the presence
of 0.5 mg/L of nitrite. If nitrification episodes are allowed to continue, very low (or zero) total
chlorine residual concentration levels may occur. Loss of chlorine residual allows an increase in
HPC bacteria and potentially increases in total coliforms and may result in a positive sample
(Cowman and Singer, 1994). Additional information on nitrification can be found in (Kirmeyer et,
al. 1995), "Nitrification Occurrence and Control in Chloraminated Water Systems."
Factors. Several possible factors have been implicated as contributing to nitrification. These factors
include low chlorine-to-ammonia ratio, long detention times, and temperatures (Cowman and Singer,
1994). Though some articles noted that low monochloramine dosages may lead to nitrification, other
research has reported nitrification occurring at monochloramine concentrations greater than 5.0 mg/L
(Cowman and Singer, 1994). Nitrifying bacteria are relatively more resistant to disinfection by
monochloramine than free chlorine (Cowman and Singer, 1994). The optimum conditions for
nitrification would be a water system with free-ammonia, a pH of 7.5 to 8.5, a water temperature of
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6. CHLORAMINES
25 to 30°C, and a dark environment. Nitrifying bacteria exhibit slow growth and have been found in
higher numbers in the sediment of distribution systems than in the biofilm (Cowman and Singer,
1994).
If the water reservoirs in the distribution system are covered, partial nitrification may occur (White,
1992). Nitrification occurred in two of MWDSC's covered reservoirs (Garvey and Orange County
reservoirs) after the changeover. Approximately 10 weeks after the changeover from chlorine to
chloramine, water quality degradation was occurring in the Garvey Reservoir. MWDSC increased
the amounts of free chlorine added to the plant effluent to maintain a 1.5 mg/L monochloramine
residual at the reservoir effluent and the chlorine to ammonia ratio was increased from 3:1 to 4:1 to
decrease the amount of excess ammonia in the water. These changes were more effective than the
flushing programs for the distribution system, which only helped temporarily.
Control Measures. Nitrification may pose a potential problem for any utility using monochloramine
as a disinfectant (Cowman and Singer, 1994). Thus, nitrification should be carefully assessed and
controlled. Nitrification may be controlled by monitoring at strategic locations throughout the
distribution system for monochloramine and dichloramine residuals (White, 1992).
The chloramine free residual stability is increased throughout the distribution system when there is
increased control of microbial contaminants and decreased bacterial concentrations in the raw water
to acceptable levels. Recommended approaches to prevent and control nitrification in the
distribution system include (Cowman and Singer, 1994):
• Decreasing the detention time;
• Increasing the pH;
• Decreasing the temperature;
• Decreasing TOC concentrations;
• Increasing chloramines residual;
• Increasing the chlorine-to-ammonia ratio; and
• Decreasing the excess ammonia concentration.
For the distribution system, the system should be evaluated to identify the low-flow or dead-end
sections. The detention times in the system should be operationally minimized (Skadsen, 1993). For
reservoirs, those with single inlet-outlet configurations should especially be carefully monitored and
operated (Skadsen, 1993).
MWDSC stresses the importance of developing nitrification strategy control measures. In particular,
a comprehensive monitoring program should be established to alert personnel to implement control
measures when required. To control nitrification, MWDSC developed a control strategy where the
reservoirs and distribution system were first sampled for nitrite levels (Skadsen, 1993). MWDSC
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6. CHLORAMINES
also decreased the detention times in reservoirs and distribution systems, especially during warmer
weather, which helped to keep nitrite levels down.
The chloramination operation was modified to add more chlorine to the reservoir inlet and increase
the chlorine-to-ammonia ratio from 3:1 to 5:1 at the plant effluent. The initial 3:1 ratio corresponded
with a 1.5 mg/L monochloramine residual and 0.2 mg/L excess ammonia. At these concentrations,
the agencies receiving the water had the flexibility of blending the chloraminated water with
chlorinated water or adding more chlorine to those sections of the system with long detention times.
Increasing the ratio further to 5:1 controlled the nitrification problem by decreasing the amount of
free ammonia in the distribution system. Operating at a 5:1 ratio requires more monitoring since an
overdose of chlorine can reduce the chloramine residual.
A survey of chloramine users in the United States was conducted in June 1991. This survey showed
that the chlorineto ammonia ratio varied from 3:1 to 7:1 (Dennis, et al, 1991). The chloramine
residual varied from 0.8 mg/L to 3 mg/L (Dennis, et al, 1991). Table 6-7 shows the results from this
survey. The agencies surveyed reported excellent results with secondary chloramine disinfection
(Dennis et al, 1991). EPA is in the process of collecting and evaluating chloramine use in the United
States as part of the Information Collection Rule (ICR), but until those data are available, the 1991
survey appears to be the most recent national survey of chloramine use.
Each year, MWDSC also added chlorine past the breakpoint to allow a free residual for 30 days. The
ideal locations for breakpoint chlorination are at the distribution reservoirs and interconnections. The
increased chlorine oxidizes any nitrite and nitrifying bacteria and eliminates the excess ammonia in
the distribution system. For larger water systems, MWDSC recommends maintaining chlorination
stations throughout the distribution system. Both fixed and mobile chlorinators may be used. Mobile
chlorinator units are self-contained and trailer-mounted with evaporators, chlorinators, generator, a
booster pump for transport water, and chlorine injectors. They are designed to draw liquid chlorine
directly from a 17-ton chlorine trailer and to inject a chlorine solution into the distribution system or
reservoir.
Since nitrifying bacteria were found in higher numbers in the sediments of the distribution system
than in the biofilm, flushing sediment from the system will help to control nitrification. The addition
of a disinfectant (i.e., free or combined chlorine) is required to remove nitrification.
At the Indiana American Water Company, the distribution system is temporarily converted back to
free chlorine for scheduled flushing (Lyn et al., 1995). Utilities should evaluate their flushing
program to avoid consumer complaints with inappropriate flushing techniques.
6.7.2.3 Taste and Odor
If the chlorine to ammonia-nitrogen ratios are between 3:1 and 5.5:1, disagreeable tastes and odors
should be evaluated at the consumer tap (White, 1992).
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6. CHLORAMINES
Fishy tastes and odors (e.g., from source waters and return washwater from the washwater treatment
system) can be controlled by a 1-hour contact time with free-chlorine residual of 2 mg/L prior to the
addition of ammonia (Dennis et al., 1991). This prechlorination eliminates the fishy taste and odor
but may increase the THM concentrations at the plant effluent.
Table 6-7. Survey of Chloramine Users in the United States
Agency
City of
Dallas, TX
City of
Denver, CO
Indianapolis,
Water Co.,
IN
Miami-Dade
Water
Authority, FL
City of
Milwaukee,
Wl
City of
Philadelphia
,PA
City of
Portland,
OR
Orleans
Parish, LA
St. Louis
Co. Water
Authority,
MO
Treatment
Capacity
730 mgd;
3 plants
600 mgd;
3 plants
176 mgd;
4 plants
300 mgd;
3 plants
305 mgd;
2 plants
530 mgd;
3 plants
225 mgd
300 mgd;
2 plants
360 mgd;
4 plants
Type of
Ammonia
Anhydrous
Aqueous
(30%)
Anhydrous
Anhydrous
Anhydrous
Aqueous
(30%)
Anhydrous
Anhydrous
Aqueous
(30%)
Chlorine:
Ammonia
Nitrogen
ratio
5:1
3:1
3:1
varies
5:1
5:1
3:1
7:1
3:1
4:1
Chloramine
Residual
(mg/L)
2.1 -2.3
0.8-1.0
1.5-2.0
2.7 - 3.0
0.8 - 0.9
2.0
1.8
2.0 - 2.5
2.5
NH3 Injection
Point
Presedimenta
-tion, Post-
filtration
Post-filtration,
prior to
chlorine
addition
Post-filtration
10 ft after
chlorine flash
mix
Post-filtration
Post-filtration
70ft
downstream
of chlorine in
conduit
Pre-filtration
Concurrent
with chlorine
at flash mix,
post-filtration
Nitrification
Control
Strategies
None
None
Increase ratio in
summer
2 weeks free
chlorine every
November
None
None
None
None
None
Source: Dennis etal., 1991.
6.7.3 Special Considerations for Chloramination Facilities
6.7.3.1 GAC Filters with Ammonia Addition
The Ann Arbor Water Treatment Plant in Michigan is a 50 mgd lime softening plant that draws its
water from the Huron River (80 to 90 percent) and ground water (10 to 20 percent). When
chloramination is applied to the river water, the chlorine is injected into the raw water line
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6. CHLORAMINES
immediately before ammonia is applied. The total chlorine feed averages 3.3 mg/L with an average
demand of 2.0 mg/L for the river water.
Evidence of nitrification occurred immediately after a change in treatment from sand to GAC
filtration. Prior to the change to GAC, the treatment plant had successfully used monochloramine as
both a primary and a final disinfectant. Nitrification was not evident. The GAC received an
application of approximately 1.3 mg/L ammonia. This input of ammonia to the filters constituted a
nutrient source that allowed nitrifying bacteria to become established and proliferate. The GAC
particles have been observed to harbor nitrifying bacteria and nitrification has been observed in GAC
beds. Higher nitrifying bacteria levels have been observed in other filter beds as compared with
source water. The GAC effluent also showed pronounced seasonal peaks in HPC bacteria from May
to July, and percent total coliforms positive from July to August. These seasonal peaks are most
probably temperature related. During periods of nitrification, GAC effluent HPC bacteria
concentration was steadily decreasing while in the distribution system, HPC bacteria were increasing.
6.7.3.2 Organic Nitrogen
Concentrations of organic nitrogen and ammonia nitrogen as low as 0.3 mg/L may interfere with the
chloramination process! The monochloramine residuals will hydrolyze with the organic nitrogen to
form organochloramines, which are nongermicidal. This reaction would take about 30 to 40 minutes.
After the monochloramine residuals disappear, free ammonia nitrogen reappears. Free ammonia
nitrogen is a powerful biological nutrient. Its presence promotes biological instability in that portion
of the distribution system. Biological instability usually results in foul tastes and odors plus dirty
and/or colored water at the consumers tap (White, 1992).
The free chlorine residual or chloramine residual method may be used to clean an area with
biological instability. Of the two methods, the free chlorine residuals method is superior (White,
1992). Free chlorine residuals restore distribution system stability quicker (i.e., a few days for free
chlorine versus weeks for chloramines), the clean-up process can be monitored, and the clean-up is
complete when the free chlorine residual concentration reaches 85 percent of the free chlorine
concentration.
Based on their conversion to chloramination experience, MWDSC recommends that utilities
employing chloramines for disinfection monitor for total organic nitrogen levels. When levels are
high, the amino acid fraction is also likely to rise. This rise may impair the chloramination
disinfection efficiency if high levels of organic nitrogen are not detected.
6.7.3.3 Mixing
Mixing at the point of application greatly affected the bactericidal efficiency of the chloramine
process. When the pH of the water is between 7 and 8.5, the reaction time between ammonia and
chlorine is practically instantaneous. If chlorine is mixed slowly into the ammoniated water, organic
matter, especially organic matter prone to bleaching with chlorine solution, may react with the
chlorine and interfere with chloramine formation (White, 1992).
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6. CHLORAMINES
6.7.3.4 Blending Waters
When chlorinated water is blended with chloraminated water, the chloramine residual will decrease
after the excess ammonia has been combined and monochloramine is converted to dichloramine and
nitrogen trichloride. The entire residual can be depleted. Therefore, it is important to know how
much chlorinated water can be blended with a particular chloraminated water stream without
significantly affecting the monochloramine residual. Blended residual curves should be developed
for each specific blend.
6.7.3.5 Corrosion
Chloramination and corrosion control can limit bacterial biofilm development in the distribution
system. If optimum corrosion of iron pipes is not controlled, the chloramination efficiency may be
impacted. Corrosion inhibitors with higher phosphate concentrations may reduce corrosion rates
(Lynetal., 1995).
6.7.3.6 Formation of Nitrogen Trichloride
If water in the distribution system tends to form nitrogen trichloride, the finished water should be
subjected to post-aeration, which readily removes nitrogen trichloride (White, 1992). Nitrogen
trichloride is also readily destroyed by sunlight (White, 1992).
6.7.3.7 Human Health and the Environment
Users of kidney dialysis equipment are the most critical group that can be impacted by chloramine
use. Chloramines can cause methemoglobinemia and adversely affect the health of kidney dialysis
patients if chloramines are not removed from the dialysate water. Chloramines can also be deadly to
fish. The residuals can damage the gill tissues, enter the red blood cells, and cause an acute blood
disorder. Chloramine residuals should be removed from the water prior to the water contacting any
fish. As such, fish hobbyists should be notified, along with pet stores and aquarium supply
establishments.
6.8 Summary
6.8.1 Advantages and Disadvantages of Chloramine Use
The following list highlights selected advantages and disadvantages of using chloramines as a
disinfection method for drinking water (Masschelein, 1992). Because of the wide variation of
system size, water quality, and dosages applied, some of these advantages and disadvantages
may not apply to a particular system.
Advantages
• Chloramines are not as reactive with organics as free chlorine in forming DBFs.
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6. CHLORAMINES
• The monochloramine residual is more stable and longer lasting than free chlorine or
chlorine dioxide, thereby providing better protection against bacterial regrowth in
systems with large storage tanks and dead end water mains. However excess ammonia in
the network may cause biofilming.
• Because chloramines do not tend to react with organic compounds, many systems will
experience less incidence of taste and odor complaints when using chloramines.
. Chloramines are inexpensive.
• Chloramines are easy to make.
Disadvantages
« The disinfecting properties of chloramines are not as strong as other disinfectants, such as
chlorine, ozone, and chlorine dioxide.
« Chloramines cannot oxidize iron, manganese, and sulfides.
. When using chloramine as the secondary disinfectant, it may be necessary to periodically
convert to free chlorine for biofilm control in the water distribution system.
. Excess ammonia in the distribution system may lead to nitrification problems, especially
in dead ends and other locations with low disinfectant residual.
• Monochloramines are less effective as disinfectants at high pH than at low pH.
. Dichloramines have treatment and operation problems.
• Chloramines must be made on-site.
6.8.2 Summary Table
Table 6-8 summarizes the considerations for the use of chloramine.
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6. CHLORAMfNES
Table 6-8. Summary of Chloramine Disinfection
Consideration
Description
Generation Chloramines are generated by the sequential addition of chlorine (hypochlorous acid) and
ammonia at a Cb to NHa ratio ranging from of 3:1 to 5:1. Either chlorine or ammonia may be added
first. Chlorine is normally added first to act as the primary disinfectant and after 10 to 30 minutes,
ammonia is added to prevent further formation of DPBs.
The most common methods of chlorine addition include gas feed using a dilution water eduction
system or direct feed of bulk hypochlorite solution (12 percent typical commercial strength).
The most common ammonia feed facilities include anhydrous ammonia fed either directly or via a
dilution water eduction system or direct feed of bulk aqua ammonia solution (20 percent typical
commercial strength).
Primary uses Monochloramine is used primarily as a secondary disinfectant to provide a residual in the
distribution system. It is used where elevated DBPFP levels in the treated water can cause high
levels of DBP formation in the distribution system if free chlorine is used as the secondary
disinfectant. Monochloramine has been found to be more effective than free chlorine in controlling
biofilms and coliform bacteria in systems with long detention times due the lower decay rate of
chloramine. Monochloramine will have much less tendency to react with organics present and
hence will form' less taste and odor causing compounds.
Inactivation efficiency At pH 7 and below, free chlorine is 200,200,50, and 2.5 times more effective in inactivating
bacteria, viruses, spores, and cysts respectively than monochloramine.
Byproduct formation Monochloramine substantially reduces the DBP formation but still forms some DBPs.
Limitations Monochloramine is increasingly being used as a secondary disinfectant to provide a residual in
distribution systems because of its lower decay rate then free chlorine and lesser tendency to form
DBPs.
Caution should be used in using monochloramine in distribution systems where water sources
using free chlorine residual are also used. High Cb to N ratios can occur where waters using
different residuals combine leading to the possible formation of taste and odor causing
dichloramine and nitrogen trichloride. In some cases the residual maybe completely removed by
the breakpoint reaction.
Point of application
Special considerations
Monochloramine is normally generated at the treatment facility with the addition of ammonia to
chlorinated water. Ammonia is normally added prior to the pumping into the distribution system. In
some cases, ammonia is added prior to the clean/veil to minimize formation of DBPs by free
chlorine residual.
Nitrification and generation of bacterial growths can occur if the Cb to N ratio is too low and
conditions exist for the growth of nitrifying bacteria. A minimum residual of 2.0 mg/L of
monochloramine has been found effective in controlling nitrification in most systems.
6.9 References
1. Aoki, T. 1989. "Continuous Flow Method For Simultaneous Determination Of Monochloramine,
Dichloramine, and Free Chlorine: Application To A Water Purification Plant." Environ. Sci.
Technol. 23:46-50.
April 1999
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6. CHLORAMINES
2. AWWA (American Water Works Association). 1991. Guidance Manual for Compliance with the
Filtration and Disinfection Requirements for Public Works Systems Using Surface Water
Sources.
3. AWWA and ASCE (American Society of Civil Engineers). 1990. Water Treatment Plant Design,
second edition. McGraw-Hill, Inc. New York, NY.
4. Bongers, L.H., T.P O'Connor, and D.T. Burton. 1977. "Bromine Chloride-An Alternative To
Chlorine For Fouling Control in Condenser Cooling Systems." EPA 600/7-77-053, Washington,
D.C.
5. Cameron, G.N., J.M. Symons, S.R. Spencer, and J.Y. Ma. 1989a. "Minimizing THM Formation
During Control of the Asiatic Clam: A Comparison of Biocides." /. AWWA. 81(10):53-62.
6. Cameron, G.N., J.M. Symons, D. Bushek, and R. Kulkarni. 1989b. "Effect of Temperature and
pH on the Toxicity of Monochloramine to the Asiatic Clam." J. AWWA. 81(10):63-71.
7. Chang, S.L. 1944. "Studies on Entamoeba histolytica 3. Destruction of Cysts of Entamoeba
histolytica by Hypochlorite Solution, Chloramines in Tap Water and Gaseous Chlorine in Tap
Water of Varying Degrees of Pollution." War Med. 5:46.
8. Chang, S.L. and G.M. Fair. 1941. "Viability and Destruction of the Cysts of Entamoeba
histolytica." J. AWWA. 33(10):1705.
9. Cowman, G.A., and P.C. Singer. 1994. "Effect of Bromide Ion on Haloacetic Acid Speciation
Resulting from Chlorination and Chloramination of Humic Extracts." Conference proceedings,
AWWA Annual Conference, New York, NY.
. jl!
10. Dennis, J.P., D.C. Rauscher, and D.A. Foust. 1991. "Practical Aspects of Implementing
Chloramines." Conference proceedings, AWWA Annual Conference, Philadelphia, PA.
11. Dorn, J. M. 1974. A Comparative Study of Disinfection on Viruses and Bacteria by
Monochloramine. Master's thesis, Univ. Cincinnati, Ohio.
12. Esposito, M.P. 1974. The Inactivation of Viruses in Water by Dichloramine. Master's thesis,
Univ. Cincinnati, Ohio.
13. Feng, T.H. 1966. "Behavior of Organic Chloramines." J. Water Pollution Control Fed.
38(4):614.
14. Fujioka, R.S., K.M. Tenno, and P.C. Loh. 1983. "Mechanism of Chloramine Inactivation of
Poliovirus: A Concern for Regulators." Water Chlorination: Environmental Impacts and Health
Affects, Vol. 4, R.L. Jolley, et al. (editor). Ann Arbor Science Publishers, Inc., Ann Arbor, MI.
EPA Guidance Manual
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:,: i Jl.: i ,i
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6. CHLORAMINES
15. Gordon, G., WJ. Cooper, R.G. Rice, and G.E. Pacey. 1992. Disinfectant Residual Measurement
Methods. Second Edition, AWWARF and AWWA.
16. Haas, C. N. and R.S. Engelbrecht. 1980. "Chlorine Dynamics During Inactivation of Coliforms,
Acid-Fast Bacteria and Yeasts." Water Res. 14:1744.
17. Hazen and Sawyer. 1992. Disinfection Alternatives For Safe Drinking Water. Van Nostrand
Reinhold, New York, NY,
18. Hoff, J.C. andE.E. Geldreich. 1981. "Comparison of the Biocidal Efficiency of Alternative
Disinfectants." /. AWWA. 73(1):40.
19. Isaac, R.A. and J.C. Morris. 1980. "Rates of Transfer of Active Chlorine Between Nitrogenous
Substances." Water Chlorination: Environmental Impact and Health Affects, Vol. 3. R.L. Jolley
(editor). Ann Arbor Science Publishers, Inc., Ann Arbor, ML
20. Jacangelo, J.G., Olivieri, V.P., and Kawata, K., 1987. "Mechanism of Inactivation of
Microorganisms by Combined Chlorine." AWWA Research Foundation, Denver, CO.
21. Jacangelo, J.G., N.L. Patania, K.M. Reagan, E.M. Aieta, S.W. Krasner, and M.J. McGuire. 1989.
"Impact of Ozonation on the Formation and Control of Disinfection Byproducts in Drinking
Water." /. AWWA.81(8):74.
22. Jensen, J., J. Johnson, J. St. Aubin, R. Christman. 1985. "Effect of Monochloramine on Isolated
Fulvic Acid." Org. Geochem. 8(1):71.
23. Johnson, J.D. 1978. "Measurement and Persistence of Chlorine Residuals." Natural Waters. In
Water Chlorination: Environmental Impact and Health Effects. R.L. Jolley (editor). Ann Arbor
Science Publishers, Inc., Ann Arbor, ML 1:37-63.
24. Kabler, P.W., et al. 1960. "Viricidal Efficiency of Disinfectants in Water." Public Health Repts.
76(7):565.
25. Kelley, S.M. and W.W. Sanderson. 1958. "The Affect of Chlorine in Water on Enteric Viruses."
Amer. Jour. Publ. Health. 48:1323.
26. Kelley, S.M. and W.W. Sanderson. 1960. "The Effect of Chlorine in Water on Enteric Viruses 2,
The Effect of Combined Chlorine on Poliomyelitis and Coxsackie Viruses." Amer. Jour. Publ.
Health. 50(1): 14.
27. Kirmeyer, G., et al. 1993. Optimizing Chloramine Treatment. AWWARF.
28. Kirmeyer, G., et al. 1995. Nitrification Occurrence and Control in Chloraminated Water Systems.
AWWARF.
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29. Krasner, S.W., MJ. McGuire, and J.J. Jacangelo. 1989. "The Occurrence of Disinfection
Byproducts in U.S. Drinking Water." 7. AWWA. 81(8):41.
30. Lyn, T.L., S.R. Lavinder, and R. Hungate. 1995. "Design Considerations for Ammoniation
Facilities." Conference proceedings, AWWA Annual Conference, Anaheim, CA.
31. Margerum, D.W., et al. 1978. "Chlorination and the Formation of N-Chloro Compounds in
Water Treatment." Organometals and Organometal-loids: Occurrence and Fate in the
Environment. R. F. Brinckman and J. M. Bellama (editors). ACS (American Cancer Society),
Washington, D.C.
32. Marks, H.C., D.B. Williams, and G.U. Glasgow. 1951. "Determination of Residual Chlorine
Compounds." /. AWWA. 43:201-207.
33. Masschelein, W.J. 1992. "Unit Processes in Drinking Water Treatment." Marcel Decker D.C.,
New York, Brussels, Hong Kong.
34. Montgomery, J.M. 1985. Water Treatment Principles and Design. John Wiley & Sons, Inc.,
New York, NY.
35. Morris, J.C. 1967. "Kinetics of Reactions Between Aqueous Chlorine and Nitrogen
Compounds." Principles and Applications of Water Chemistry. S.D. Faust and J.V. Hunter
(editor). John Wiley & Sons, New York, NY.
36. NRC (National Research Council). 1980. Drinking Water and Health, Vol. 2. National Academy
Press, Washington, D.C.
37. Norton, C.D. and M.W. LeChevallier. 1997. "Chloramination: Its Effect on Distribution System
Water Quality." J.A WWA. 89(7):66.
38. Olivieri, V.P., et al. 1980. "Reaction of Chlorine and Chloramines with Nucleic Acids Under
Disinfection Conditions." Water Chlorination: Environmental Impact and Health Affects, Vol. 3.
R.J. Jolley (editor), Ann Arbor Science Publishers, Inc., Ann Arbor, MI.
39. Palin, A. 1950. 1950. "A Study of the Chloro Derivatives of Ammonia." Water and Water
Engineering. 54:248-258.
40. Rice, R. and M. Gomez-Taylor. 1986. "Occurrence of By-Products of Strong Oxidants Reating
With Drinking Water Contaminants - Scope of the Problem." Environ. Health Perspectives.
69:31.
41. Singer, P.C. 1993. "Trihalomethanes and Other Byproducts Formed From the Chlorination of
Drinking Water." National Academy of Engineering Symposium on Environmental Regulation:
Accommodating Changes in Scientific, Technical, or Economic Information, Washington, D.C.
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6. CHLORAMINES
42. Skadsen, J. 1993. "Nitrification in a Distribution System." 7. AWWA. 95-103.
43. Smith, M.E., Cowman, G.A., Singer, P.C. 1993. "The Impact of Ozonation and Coagulation on
DBF Formation in Raw Waters." Conference proceedings, AWWA Annual Conference, San
Antonio, TX.
44. Speed, M.A., et al. 1987. Treatment Alternatives for Controlling Chlorinated Organic
Contaminants in Drinking Water. EPA/60012-87/011, Washington, D.C.
45. Standard Methods. 7995. Standard Methods for the Examination of Water and Wastewater,
ninteenth edition, Franson, M.H., Eaton, A.D., Clesceri, L.S., and Greenberg, A.E., (editors).
APHA (American Public Health Association), AWWA, and Water Environment Federation,
Washington D.C.
46. Stringer, R. and C.W. Kruse. 1970. "Amoebic Cysticidal Properties of Halogens." Conference
proceedings, National Specialty Conference on Disinfection, ASCE, New York.
47. Sugam, R. 1983. "Chlorine Analysis: Perspectives For Compliance Monitoring." Water
Chlorination, Environmental Impact and Health Effects. R.L. Jolley, et al. (editor). Ann Arbor
Science Publishers, Ann Arbor, MI.
48. Valentine, R.L. et al. 1998. "Chloramine Decomposition in Distribution System and Model
Waters." AWWA, Denver, CO.
49. Wattie, E. and C.T. Butterfield. 1944. "Relative Resistance of Escherichia coli and Eberthella
typhosa to Chlorine and Chloramines." Public Health Repts. 59:1661.
50. Weber, G.R. and M. Levine. 1944. "Factors Affecting the Germicidal Efficiency of Chlorine and
Chloramine." Amer. J. Public Health: 32:719.
51. Weil, I. and J.C. Morris. 1949. "Kinetic Studies on the Chloramines. The Rates of Formation of
Monochloramine, N-Chlormethylamine and N-Chlordimethylamine." J. Amer. Chem. Soc.
71:1664.
52. White, G.C. 1992. Handbook Of Chlorination and Alternative Disinfectants. Volume 3. Van
Nostrand Reinhold Co., New York, NY.
53. Whittle, G.P. and Lapteff, A., Jr. 1974. "New Analytical Techniques For The Study Of Water
Disinfection." Chemistry of Water Supply, Treatment, and Distribution. A.J. Rubin (editor) Ann
Arbor Science Publishers, Inc., Ann Arbor, MI.
54. Wolfe, R.L., N.R. Ward, and B.H. Olson. 1984. "Inorganic Chloramines as Drinking Water
Disinfectant: A Review."/ AWWA. 76(5):74-88.
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7. PEROXONE (OZONE/HYDROGEN
PEROXIDE)
Advanced oxidation processes generate highly reactive hydroxyl free radicals to oxidize various
compounds in the water. As discussed in Chapter 3, hydroxyl radicals are produced during the
spontaneous decomposition of ozone. By accelerating the ozone decomposition rate, the hydroxyl
radical concentration is elevated, which increases the oxidation rate. This procedure increases the
contribution of indirect oxidation over direct ozone oxidation as discussed in Chapter 3.
Several methods have been used to increase ozone decomposition and produce high concentrations
of hydroxyl radicals. One of the most common of these processes involves adding hydrogen
peroxide to ozonated water, a process commonly referred to as peroxone.
The Metropolitan Water District of Southern California (MWDSC) conducted extensive research
into the use of peroxone to control organics and oxidize taste and odor compounds (e.g., geosmin
and 2-methylisoborneol [MIB]) while providing sufficient levels of molecular ozone to guarantee CT
values and primary disinfection. While this chapter focuses on peroxone as a disinfectant, similar
results are expected from other advanced oxidation processes such as ozone plus UV, ozone at high
pH, hydrogen peroxide plus UV, and other combinations.
A key issue with the use of peroxone as a disinfection process is that the process does not provide a
measurable disinfectant residual. It is therefore not possible to calculate CT similar to the use of
other disinfectants. While no credit can be given for hydroxyl free radicals because it cannot be
measured directly, some credit may be considered for any detected ozone in peroxone systems.
Peroxone does provide pathogen inactivation, as discussed in this chapter, but equivalent CT values
or methods of calculating equipment CT credits have not been established at the date of publication
of this guidance document.
7.1 Peroxone Chemistry
The ozone decomposition cycle is similar to that discussed in Chapter 3. However, the added
hydrogen peroxide or ultraviolet radiation accelerates the decomposition of ozone and increases the
hydroxyl radical concentration. By adding hydrogen peroxide, the net production of hydroxyl free
radicals is 1.0 mole hydroxyl radical per mole ozone.
Similar to the discussion of ozone in Chapter 3, oxidation in the peroxone occurs due to two
reactions (Hoigne and Bader, 1978):
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
• Direct oxidation of compounds by aqueous ozone (O3(aq)); and
• Oxidation of compounds by hydroxyl radicals produced by the decomposition of ozone.
The two oxidation reactions compete for substrate (i.e., compounds to oxidize). The ratio of direct
oxidation with molecular ozone is relatively slow (lO'Mo'MT'sec"1) compared to hydroxyl radical
oxidation (1012-10l4M"1sec~1), but the concentration of ozone is relatively high. On the other hand,
the hydroxyl radical reactions are very fast, but the concentration of hydroxyl radicals under normal
ozonation conditions is relatively small.
A key difference between the ozone and peroxone processes is that the ozone process relies heavily
on the direct oxidation of aqueous ozone while peroxone relies primarily on oxidation with hydroxyl
radical. In the peroxone process, the ozone residual is short lived because the added peroxide greatly
accelerates the ozone decomposition. However, the increased oxidation achieved by the hydroxyl
radical greatly outweighs the reduction in direct ozone oxidation because the hydroxyl radical is
much more reactive. The net result is that oxidation is more reactive and much faster in the
peroxone process compared to the ozone molecular process. However, because an ozone residual is
required for determining disinfection CT credit, peroxone may not be appropriate as a pre-
disinfectant.
The peroxone process utilizes oxidation by hydroxyl radicals. The oxidation potential of the
hydroxyl radical and ozone are as follows:
O. + 2/T +2e~-+O2 +H2O £° = +2.07V
O3 + H2O + 2e- -*O2+ 2OH~ £° = +1.24V
In addition to having an oxidation potential of hydroxyl radical higher than ozone, the hydroxyl
radical is also much more reactive approaching the diffusion control rates for solutes such as
aromatic hydrocarbons, unsaturated compounds, aliphatic alcohols, and formic acid (Hoigne and
Bader, 1976).
7.1 .1 Oxidation Reactions
Because the radical oxidation is much more effective than direct oxidation with ozone, it has been
used extensively to treat difficult to oxidize organics such as taste and odor compounds and
chlorinated organics (e.g., geosmin, MIB, phenolic compounds, trichloroethylene [TCE], and
perchloroethylene [PCE]).
Neither ozone nor peroxone significantly destroys TOC. Peroxone will oxidize the saturated
organics and produce byproducts similar to those found in ozonation; namely, aldehydes, ketones,
EPA Guidance Manual 7-2 April 1999
Alternative Disinfectants and Oxidants
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
peroxides, bromate ion, and biodegradable organics (MWDSC and JMM, 1992). However, with
peroxone, the biodegradibility of the water (not the organic compounds) increases, rendering "a
portion of the TOC" amenable to removal in biologically active filters.
Peroxone has found a niche in oxidizing difficult-to-treat organics, such as taste and odor compounds
including geosmin and MIB (Pereira et al., 1996; Ferguson et al., 1990). In addition, peroxone and
other advanced oxidation processes have been shown to be effective in oxidizing halogenated
compounds such as 1,1-dichloropropene, trichloroethylene, 1-chloropentahe, and 1,2-dichloroethane
(Masten and Hoigne, 1992; Aieta et al., 1988; Glaze and Kang, 1988). Hydroxyl radicals will react
with all these compounds plus refractory aliphatics such as alcohols and short-chain acids (Chutny
and Kucera, 1974).
The optimum peroxide:ozone dose ratio to maximize hydroxyl radicals' reaction rate can be
determined for a specific oxidation application. For instance, the optimum peroxide:ozone dose ratio
for TCE and PCE oxidation in a ground water was determined to be 0.5 by weight (Glaze and Kang,
1988). Tests showed that TCE required lower ozone dosages for the same percentage removal
compared to PCE.
LADWP conducted pilot studies and operated a 2,000 gpm full scale AOP demonstration plant in
1995. The peroxide:ozone dose ratio used was 0.5 to 0.6. Ground water containing up to 447 mg/L
TCE and 163 mg/L PCE was treated to below the respective MCLs. However, bromate ion was
formed in excess of the 0.010 mg/L MCL (Karimi et al., 1997).
7.1.2 Reactions with Other Water Quality Parameters
As with ozone alone, pH and bicarbonate alkalinity play a major role in peroxone effectiveness
(Glaze and Kang, 1988). This role is primarily related to bicarbonate and carbonate competition for
hydroxyl radical at high alkalinity and carbonate competition for hydroxyl radical at high pH levels.
Also, excessive peroxide can also limit the formation of the hydroxyl radical and reduce the
effectiveness of peroxone.
Turbidity alone does not appear to play a role in peroxone effectiveness nor does peroxone appear to
remove turbidity. Tobiason et al. (1992) studied the impact of pre-oxidation on filtration and
concluded that the pre-oxidation did not improve effluent turbidities, but did appear to increase filter
run times because of lower head loss or delayed turbidity breakthrough. Filter effluent turbidities
were similar for no-oxidant and pre-oxidant trains.
7.1.3 Comparison between Ozone and Peroxone
The key difference between ozone and peroxone is in the primary oxidation mode; that is, direct
oxidation or hydroxyl radical oxidation. The reactivities of these compounds create a different effect
in the reactions with water constituents and, thus, disinfection effectiveness. Table 7-1 summarizes
Anril 1QQ9 " " 7-3 EPA Guidance Manual
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7, PEROXONE (OZONE / HYDROGEN PEROXIDE)
the key differences between ozone and peroxone as they relate to their application in drinking water
treatment.
Table 7-1. Comparison between Ozone and Peroxone Oxidation
Process
Ozone
Peroxone
Ozone decomposition
rate
"Normal" decomposition producing
hydroxyl radical as an intermediate
product
Accelerated ozone decomposition
increases the hydroxyl radical
concentration above that of ozone
alone.
Ozone residual
Oxidation path
Ability to oxidize iron
and manganese
Ability to oxidize taste
and odor compounds
Ability to oxidize
chlorinated organics
Disinfection ability
5-10 minutes
Usually direct aqueous molecular
ozone oxidation
Excellent
Variable
Poor
Excellent
Very short lived due to rapid reaction.
Primarily hydroxyl radical oxidation.
Less effective.
Good, hydroxyl radical more reactive
than ozone.
Good, hydroxyl radical more reactive
than ozone.
Good, but systems can only receive
Ability to detect
residual for
disinfection monitoring
Good
CT credit if they have a measurable
ozone residual.
Poor. Cannot calculate CT value for
disinfection credit.
7.2 Generation
The peroxone process requires an ozone generation system as described in Chapter 3 and a hydrogen
peroxide feed system. The process involves two essential steps: ozone dissolution and hydrogen
peroxide addition. Hydrogen peroxide can be added after ozone (thus allowing ozone oxidation and
disinfection to occur first) or before ozone (i.e., using peroxide as a pre-oxidant, followed by
hydroxyl radical reactions) or simultaneously. Addition of hydrogen peroxide following ozone is the
best way to operate, however a system cannot obtain a CT credit unless the ozone residual is
sufficiently high.
There are two major effects from the coupling of ozone with hydrogen peroxide (Duguet et al.,
1985):
• Oxidation efficiency is increased by conversion of ozone molecules to hydroxyl radicals; and
EPA Guidance Manual
Alternative Disinfectants and Oxidants
7-4
April 1999
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
• Ozone transfer from the gas phase to the liquid is improved due to an increase in ozone
reaction rates.
The most efficient operation is to add ozone first to obtain CT disinfection credit, followed by
peroxide for hydroxyl radical oxidation.
Ozonation can be described as occurring in two stages. In the first stage, ozone rapidly destroys the
initial oxidant demand present, thereby enhancing the ozone transfer rate into solution from the gas
phase. Addition of hydroxyl free radicals to the first stage should be minimized since the hydrogen
peroxide competes with ozone-reactive molecules (i.e., initial demand) for the ozone present. In the
second stage, organic matter is oxidized, taking place much slower than in the first stage. Adding
hydrogen peroxide during the second stage makes it possible to raise the overall oxidation efficiency,
since the reaction of hydrogen peroxide with ozone produces hydroxyl radicals enhancing chemical
reaction rates. In practice, the addition of hydrogen peroxide to the second stage of ozonation can be
achieved by injecting the hydrogen peroxide into the second chamber of an ozone contactor (Duguet
et al., 1985). The most efficient operation is to use ozone first to obtain CT credit and peroxone
second for micropollutant destruction.
Energy consumption of the peroxone process includes that for ozone generation and application, plus
for metering pumps to feed peroxide. The peroxide addition step does not require any more training
from an operator than any other liquid chemical feed system. Systems should be checked daily for
proper operation and for leaks. Storage volumes should also be checked daily to ensure sufficient
peroxide is on hand, and to monitor usage.
7.3 Primary Uses and Points of Application
Peroxone is used for oxidation of taste and odor compounds, and oxidation of synthetic organic
compounds. Peroxone is also used for the destruction of herbicides (e.g., atrazine), pesticides, and
VOCs. Peroxone is applied at points similar to ozone for oxidation. Addition of ozone first and
hydrogen peroxide second is the better way to operate. Alternatively, hydrogen peroxide can be
added upstream of ozone.
7.3.1 Primary Uses
7.3.1.1 Taste and Odor Compound Oxidation
Peroxone is used to remove taste and odor causing compounds because many of these compounds
are very resistant to oxidation, even ozone-oxidation. More recently, significant attention has been
given to tastes and odors from specific compounds such as geosmin, 2-methyliosborneol (MIB), and
chlorinated compounds. Studies at MWDSC demonstrated the effectiveness of peroxone to remove
geosmin and MIB during water treatment (Ferguson et al., 1990; Ferguson et al., 1991; Huck et al.,
1995).
April 1999 7-5 EPA Guidance Manual
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7. PEROXONE (OZONE/HYDROGEN PEROXIDE)
7.3.1.2 Synthetic Organic Compound Oxidation
Peroxone and other advanced oxidation processes have been shown to be effective in oxidizing
halogenated compounds such as l,l-dichloropropene (DCPE), TCE, 1-chloropentane (CPA), and
1,2-dichloroethane (DCA) (Masten and Hoigne, 1992; Aieta et al., 1988; Glaze and Kang, 1988).
The hydroxyl radicals formed react with all these compounds plus refractory aliphatics, such as
alcohols and short-chain acids (Chutny and Kucera, 1974).
7.3.2 Points of Application
The peroxone process is applied at points similar to ozone for oxidation as discussed in Chapter 3.
Importantly, peroxone addition should be after settling and prior to biological filtration. It is
important to add hydrogen peroxide after the initial ozone demand is consumed to avoid hydroxyl
free radical competition with the initial ozone demanding constituents.
7.3.2.1 Impact on Other Treatment Processes
Peroxide addition impacts other processes at the water treatment facility. These impacts include:
• The use of hydroxyl free radicals generates BDOC, which can cause biological growth in
distribution systems if not reduced during biologically active filtration. When peroxide
addition is placed before filters, it impacts the filters by increasing biological growths and
increasing backwash frequency (depending on the level on BDOC produced).
* Hydroxyl free radicals are strong oxidants that interfere with addition of other oxidants, such
as chlorine, until the ozone residual is quenched.
* The oxidation of iron and manganese by hydroxyl free radicals generates insoluble oxides
that should be removed by sedimentation or filtration. This also may impact the filters by
increasing the load on the filters and increasing backwash frequency.
The reader is referred to the Microbial and Disinfection Byproduct Simultaneous Compliance
Guidance Document (currently in production) for additional information regarding the interaction
between oxidants and other treatment processes.
7.4 Pathogen Inactivation
Both peroxone and other advanced oxidation processes have been proven to be equal or more
effective than ozone for pathogen inactivation. Disinfection credits are typically described in terms
of CT requirements. Because peroxone leaves no measurable, sustainable residual, calculation of an
equivalent CT for disinfection credit is not possible unless there is measurable ozone residual.
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
7A.I Inactivation Mechanism
Experiments have .indicated that long contact times and high concentrations of hydrogen peroxide
are required for bacteria and virus inactivation (Lund, 1963; Yoshe-Purer and Eylan, 1968; Mentel
and Schmidt, 1973). Achieving a 99 percent inactivation of poliovirus required either a hydrogen
peroxide dose of 3,000 mg/L for 360 minutes or 15,000 mg/L for 24 minutes. Based on these results,
when the combination of ozone and hydrogen peroxide is used, the primary cause for pathogen
inactivation is attributed to ozone, specifically the mechanisms associated with the oxidation of
pathogens by direct ozone reaction and hydroxyl radicals.
As described in Chapter 3, the mode of action of ozone on microorganisms is poorly understood.
Some studies on bacteria suggest that ozone alters proteins and unsaturated bonds of fatty acids in
the cell membrane, leading to cell lysis (Scott and Lesher, 1963; Pryor et al., 1983), while other
studies suggest that ozone may affect deoxyribonucleic acid (DNA) in the cell (Hamelin and Chung,
1974; Ohlrogge and Kernan, 1983; Ishizaki et al., 1987). Virus inactivation was reported to be
related to the attack of the protein capsid by ozone (Riesser et al., 1977). Little information was
found discussing the mode of action of ozone on protozoan oocysts. However, a few researchers
have suggested that ozone causes the oocyst density to decrease and alters the oocyst structure
(Wickramanayake, 1984; Wallisetal., 1990).
The debate continues regarding the primary mode of action for hydroxyl free radicals. Some
researchers believe that ozone disinfection is a result of direct ozone reaction (Hoigne and Bader,
1975; Hoigne and Bader, 1978), while others believe that the hydroxyl radical mechanism for
disinfection is the most important mechanism (Dahi, 1976; Bancroft et al., 1984). Studies using
ozone-hydrogen peroxide have shown that disinfection of E. coli is less effective as the peroxide to
ozone ratio increases to above approximately 0.2 mg/mg (Wolfe et al., 1989a; Wolfe et al., 1989b).
The decrease in disinfection was believed to be cause by lower ozone residuals associated with
higher peroxide to ozone ratios, which indicates that direct ozone reaction is an important
mechanism for pathogen inactivation.
7.4.2 Environmental Effects
Although the chemistry of the peroxone process is still not completely understood, optimal
production of the hydroxyl radical appears to depend on the pH, ozone concentration, ratio of
hydrogen peroxide to ozone, contact time, and water composition (Glaze et al., 1987).
7.4.2.1 Competing Chemical Reactions
One disadvantage of the peroxone process is that it involves radical intermediates that are subject to
interference from substances that react with hydroxyl radical, decreasing the effectiveness of the
process. Alkalinity, bicarbonate, and pH play a major role in the effectiveness of hydroxyl free
radicals. This effect is primarily related to bicarbonate competition for hydroxyl radical at high
alkalinity and carbonate competition for hydroxyl radical at pH levels higher than 10.3 (see Chapter
April 1999~~ 7-7~~ EPA Guidance Manual
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
3). Lowering the alkalinity prior to the application of the peroxone process may be necessary for
water that has a high bicarbonate level. In addition to carbonate and bicarbonate, organic
constituents of humic substances have also been found to react with the hydroxide radical (Glaze,
1986).
7.4.2.2 Ratio of Hydrogen Peroxide and Ozone
A study conducted at MWDSC indicated that the performance of peroxone is greatly dependent upon
the peroxide:ozone ratio (Wolfe et al., 1989b). Results from previous studies at MWDSC suggested
that the optimal ratio for disinfection was less than or equal to 0.3. One of the primary objectives of
the 1989 study was to optimize further the process for disinfection by altering peroxide:ozone ratios
and contact times. Results from the study indicated that peroxone at a 0.2 ratio of peroxide:ozone
was comparable to ozone for disinfection of indicator organisms and Giardia muris cysts, and that at
higher ratios, disinfection decreased because ozone decreased.
7.4.3 Disinfection Efficacy and Pathogen Inactivation
Recent studies have indicated that the disinfection effectiveness of peroxone and ozone are
comparable (Wolfe et al., 1989b; Ferguson et al., 1990; Scott et al., 1992). A study conducted by
Ferguson et al. (1990) compared the pathogen inactivation capability of peroxone and ozone using
MS-2 and f2 coliphages as well as E. coli. and heterotrophic plate count (HPC) bacteria. The f2 and
MS-2 coliphages were comparable in their resistance to ozone and peroxone. No differences in the
amounts of MS-2 or f2 inactivation were apparent when the peroxide:ozone ratio was varied from 0
to 0.3. Results of the E. coli. and HPC studies showed that peroxone and ozone also had comparable
results in regards to bacteria inactivation.
Table 7-2 lists CT values derived for inactivation of Giardia muris cysts by ozone and peroxone
from another study conducted by MWDSC. The contact times used for calculating the CT values
were based on 10% and 50% breakthrough of tracer compounds in the contactor. Ozone
concentrations used for CT were based on the ozone residual and half of the residual and dose. The
results of this study suggest that peroxone is slightly more potent than ozone based on the fact that
CT values for ozone were greater than for peroxone. However, because ozone decomposes more
rapidly in the presence of hydrogen peroxide, higher ozone dosages may be necessary with peroxone
to achieve comparable residuals. Moreover, the use of ozone residuals to calculate CT products for
peroxone may not take into account other oxidizing species that may have disinfectant capabilities.
Table 7-2. Calculated CT Values (mg-min/L) for the Inactivation of Giardia muris
Inactivation Ozone Ozone Peroxone" Peroxone
CTa CTa CTa CTa
90%" ~~~~^ ^g 2.8 : T2 " 2~lf
99% 3.4 5A 2.6 5.2
Data obtained from Wolfe et al., 1989b. Results at 14°C. -—-—_—__
EPA Guidance Manual 7-8 April 1999
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7. PEROXONE (OZONE/ HYDROGEN PEROXIDE)
a Ci, ozone residual; C2 (ozone dose + ozone residual)/2; Ti and T2 time (in minutes) to reach 10 percent and 50 percent
breakthrough, respectively
b The H2O2/O3 ratio for all results was 0.2.
7.5 Disinfection Byproducts
The principal byproducts associated with peroxone are expected to be similar to those for ozonation
and are listed in Table 3-9. Additional DBFs could form from reactions with hydroxyl radicals.
Peroxone does not form halogenated DBFs when participating in oxidation/reduction reactions with
NOM. However, if bromide ion is present in raw water halogenated DBFs may be formed. Similar
to ozone, the principal benefit of using peroxone for controlling THM formation appears to be that it
eliminates the need for pre-chlorination and allows lower doses of free chlorine or chloramines to be
applied later in the process train after precursors have been removed by coagulation, sedimentation,
and/or filtration and at lower doses. But peroxone does not reduce the DBPFP.
Based upon studies and findings involving peroxone, there is no beneficial lowering of THMs as
long as free chlorine is utilized as a secondary disinfectant, unless the application of peroxone allows
chlorine to be applied later in the process train to water containing reduced precursor concentrations.
The MWDSC study found that the use of peroxone/chlorine resulted in THM concentrations 10 to 38
percent greater than the use of ozone/chorine. However, the THM concentrations of waters
disinfected with peroxone/chloramines and ozone/chloramines were similar (Ferguson et al., 1990).
The use of peroxone as a primary disinfectant and chloramines as a secondary disinfectant can
successfully control halogenated DBF formation if bromide ion is not present and adequate CT credit
can be established. As with ozone, bromate ion formation is a potential concern with source waters
containing bromide ions. The oxidation reaction of bromide ion (Br~) to hypobromite ion (BrO~) and
bromite ion (BrO2") and subsequently to bromate ion (BrO3~) occurs due to direct reaction with
ozone, intermediate reactions can also occur through hydroxide radical mediated mechanisms if
bromide is not present and adequate CT credit can be established (Pereira et al., 1996).
In general, peroxone produces more bromate ion than ozone when similar ozone residuals (CT
credits) are achieved (Krasner et al., 1993). On the other hand, when the ozone dosage is kept
constant, peroxone has tended to produce comparable amounts of bromate ion as ozone. Although
peroxone produces hydroxyl radicals that can increase bromate ion formation, hydrogen peroxide
may also reduce the hypobromite ion (produced initially during the ozonation of bromide) back to
bromide ion.
A study by MWDSC evaluated the effectiveness of peroxone to control taste and odor, DBFs, and
microorganisms (Ferguson et al., 1990). In attempting to optimize the hydrogen peroxide to ozone
ratio (HaO^rOs) and the contact time for the source water, the study found pre-oxidation of source
waters followed by secondary disinfection with chloramines was an effective strategy for controlling
concentrations of THMs and other DBFs. The study found that the two source waters disinfected
with peroxone, with free chlorine as the secondary disinfectant, resulted in THM concentrations
April 1999 7-9 EPA Guidance Manual
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
ranging from 67 to 160 Hg/L. Conversely, using chloramines as a secondary disinfectant resulted in
THM concentrations consistently below 3.5 jig/L (Ferguson et al., 1990).
However, if a short free chlorine contact time is applied after biological filtration and before
ammonia addition to inactivate heterotrophic plate count bacteria in the effluent of the biologically
active filter—THMs and other DBFs will be formed at higher concentrations than from post
chloramination alone. Depending on the TOC and bromide ion concentration of the water, as well as
the pH of chlorination, temperature, and reaction time, between 2 and 28 |ag/L of TTHMs have been
formed in experiments conducted at MWDSC (personal communication, 1998).
Bromate formation in conventional ozonation, and advanced oxidation processes combining ozone
and hydrogen peroxide, were recently investigated at five water treatment plants in France. The
source water bromide ion concentrations ranged from 35 to 130 p.g/L. Bromate ion formation during
the ozonation step varied from less than 2 to 42 jog/L. In general, advanced oxidation results in
greater bromate ion concentrations when compared with conventional ozonation, provided the same
ozone residual is maintained for both processes. However, lower concentrations of bromate ion
result if the ozone dosage is kept constant between the two processes and the hydrogen peroxide
dosage is increased (von Gunten et al., 1996). To reduce bromate ion formation potential, the
proposed ozone contactor at Stone Canyon Filtration Plant includes three ozone application points
instead of two (Stolarik and Christie, 1997). Thus, when peroxone is used for obtaining CT credit,
more^bromate ion may form than during ozonation. However, if peroxone is only used for
micropollutant destruction, less bromate ion may form than when ozone is used.
7.6 Status of Analytical Methods
Hydrogen peroxide in solution reacts with ozone to ultimately form water and oxygen. Consequently,
the simultaneous presence of both oxidants is accepted as being only transient (Masschelein et al.,
1977), Chapter 3 summarizes ozone analytical methods that can be used for ozone/hydrogen
peroxide disinfectant residual monitoring. This section will present the status of analytical methods
for hydrogen peroxide only.
7.6.1 Monitoring of Hydrogen Peroxide
Standard Methods (1995) does not list procedures for measuring hydrogen peroxide. Gordon et al.
(1992) list several methods for hydrogen peroxide analysis including:
• Titration methods;
• Colorimetric methods; and
• Horseradish peroxidase methods.
EPA Guidance Manual 7-10 April 1999
Alternative Disinfectants and Oxidants
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
Table 7-3 shows the working range, expected accuracy and precision, operator skill level required,
interferences and current status for hydrogen peroxide analysis.
7.6.1.1 Titration Methods
Two titration methods are available for the analysis of hydrogen peroxide; namely, iodometric and
permanganate. Precautions for the iodometric titration include the volatility of iodine, interferences
by metals such as iron, copper, nickel, and chromium, and fading titration end points (Gordon et al.,
1992). Organic and inorganic substances that react with permanganate will interfere with the
permanganate titration.
Titration of hydrogen peroxide with permanganate or iodide ion is not sufficiently sensitive for
determining residual concentrations (Masschelein et al., 1977).
7.6.1.2 Colorimetric Methods
The most widespread method for the colorimetric determination of hydrogen peroxide is that based
on the oxidation of a Titanium (IV) salt (Masschelein et al., 1977). A yellow complex is formed and
measured by absorption at 410 nm. On a qualitative basis, ozone and persulfates do not produce the
same colored complex.
The oxidation of the leuco base of phenolphthalein is used as a qualitative test for hydrogen peroxide
(Dukes and Hydier, 1964). Sensitivity and precision of the method is sufficient in the range between
5 and 100 (J,g/L. This low working analytical range makes this method impractical for measuring
hydrogen peroxide residual levels. Also, the instability of the color obtained makes the method less
suitable for manual use. No interference data are available, but it is expected that other oxidants
would interfere (Gordon et al., 1992).
The oxidation of cobalt (II) and bicarbonate in the presence of hydrogen peroxide produces a
carbonato-cobaltate (III) complex (Masschelein et al., 1977). This complex has absorption bands at
260, 440, and 635 nm. The 260 nm band has been used for the measurement of hydrogen peroxide.
A detection limit of 0.01 mg/L has been reported (Masschelein et al., 1977). Optical interferences
are caused by 100 mg/L nitrate and 1 mg/L chlorite ions. Other oxidizing agents do interfere with
this method as will any compound with an absorption at 260 nm (Gordon et al., 1992).
April 1999 7-11 EpA Guidance Manual
Alternative Disinfectants and Oxidants
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7. PEROXONE (OZONE/ HYDROGEN PEROXIDE)
7.6.1.3 Horseradish Peroxidase Methods
Several methods incorporate the chemical reactions between peroxidase and hydrogen peroxide.
Horseradish derived peroxidase (HRP) is used most often. The scopoletin procedure is one of the
more widely accepted fluorescent methods of low levels of hydrogen peroxide utilizing HRP
(Gordon et al., 1992). Again, no information is available on potential interference.
7.6.1.4 Summary
In general, the analytical procedures for hydrogen peroxide in drinking water are impacted by other
oxidizing species such as ozone. Three of the methods are currently used, but not recommended for
disinfectant residual measurement (Table 7-3). The scopoletin HRP method is the most promising,
although additional study of potential interferences is required (Gordon et al., 1992).
7.7 Operational Considerations
Peroxide is a strong oxidant and contact with personnel should be avoided. Secondary containment
should be provided for storage tanks to contain any spills. Dual containment piping should be
considered to minimize the risk of exposure to plant personnel. Storage containers may explode in
the presence of extreme heat or fire.
7.7.1 Process Considerations
The impacts of the peroxone process are similar to those described for ozone in Chapter 3. Because
an additional oxidant is added to the water, the tendency to transform organic carbon compounds to a
more biodegradable form may be increased with the addition of peroxide.
7.7.2 Space Requirements
The metering pumps used to add peroxide should be housed with adequate space around each pump
for maintenance access. These pumps are generally not very large, so space requirements are not
significant.
The storage area can range from small where peroxide is obtained in drums, to large tank farms if
plant flow is great. As mentioned previously, secondary containment should be provided. Peroxide
has a lower freezing point than water. Housing or heat tracing should be provided for storage tanks
and exterior piping if extended periods with temperatures below freezing are anticipated.
7.7.3 Materials
Peroxide can be stored in polyethylene drums or tanks. The specific gravity is 1.39 for 50 percent
peroxide, which should be considered in the design of the tank walls. Acceptable pipe materials for
peroxide include 316 stainless steel, polyethylene, CPVC, and Teflon. Gaskets should be Teflon
April19g9 7-13 EPA Guidance Manual
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7. PEROXONE (OZONE/HYDROGEN PEROXIDE)
because natural rubber, Hypalon and EPDM are not resistant to hydrogen peroxide. Metering pumps
heads should be constructed of peroxide resistant materials.
Hydrogen peroxide is purchased from chemical suppliers and is commercially available in 35, 50,
and 75 percent strengths. Peroxide is supplied in drums or in bulk by tankcar. Price depends on
strength and quantity.
Peroxide can be stored pnsite, but deteriorates gradually over time even when stored correctly.
Pqroxide deteriorates rapidly if contaminated and with heat or exposure to certain materials.
Peroxide is added to the water with metering pumps to accurately control dose. Pumps should be
designed to prevent potential air binding of peroxide off-gas. Multiple pumps should be provided for
redundancy. As with any chemical added to water, adequate mixing should be provided.
7.8 Summary
7.8.1 Advantages and Disadvantages of Peroxone Use
(Ozone/Hydrogen Peroxide)
The following list highlights selected advantages and disadvantages of using peroxone as a
disinfection method for drinking water. Because of the wide variation of system size, water quality,
and dosages applied, some of these advantages and disadvantages may not apply to a particular
system.
Advantages
* Oxidation is more reactive and much faster in the peroxone process compared to the ozone
molecular process.
• Peroxone is effective in oxidizing difficult-to-treat organics, such as taste and odor compounds.
:•,'•• • , ... • "II ''
• Peroxone processes have been shown to be effective in oxidizing halogenated compounds.
• The tendency to transform organic carbon compounds to a more biodegradable form may be
1 1 ,i!': ' ''',', : ' i • „
increased with the addition of peroxide.
• Pumps used to house peroxide are not very large; so space requirements are not significant.
Disadvantages
• Peroxide is a strong oxidant and contact with personnel is extremely dangerous.
• Peroxide can be stored onsite, but deteriorates gradually even when stored correctly.
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7. PEROXONE (OZONE / HYDROGEN PEROX/DE)
• Peroxone as a disinfection process does not provide a measurable disinfectant residual. It is
therefore not possible to calculate CT similar to the use of other disinfectants.
• Peroxone's ability to oxidize iron and manganese is less effective than that of ozone.
7.8.2 Summary Table
Table 7-4 summarizes the considerations relative to peroxone disinfection considerations.
Table 7-4. Summary of Peroxone Disinfection Consideration
Consideration
Inactivation efficiency
Byproduct formation
Limitations
Point of application
Special considerations
Description
Generation Because of its instability, ozone must be generated at the point-of-use.
Hydrogen peroxide is purchased from chemical suppliers. Hydrogen
peroxide can be stored onsite, but is subject to deterioration.
Primary uses Primary use includes chemical oxidation. As an oxidizing agent,
peroxone can be used to remove SOC pollutants and increase the
biodegradability of organic compounds. Peroxone is an effective
disinfectant but its CT credit has not been established. It is highly
reactive and does not maintain an appreciable residual for CT credit
calculations. Peroxone may be difficult to use for disinfection because
it is highly reactive and does not maintain an appreciable ozone
residual level.
Peroxone is one of the most potent and effective germicides used in
water treatment. It is slightly more effective than ozone against
bacteria, viruses, and protozoan cysts.
Peroxone itself does not form halogenated DBPs; however, if bromide
is present in the raw water or if chlorine is added as a secondary
disinfectant, halogenated DBPs including bromate may be formed.
Other byproducts include organic acids and aldehydes.
Ideally, ozone should be used as a primary disinfectant prior to
peroxone treatment.
For disinfection, peroxone addition should be after ozonation. Ozone
contact should precede addition of hydrogen peroxide. For oxidation,
peroxone can be added prior to coagulation/sedimentation or filtration
depending on the constituents to be oxidized.
Ozone generation is a relatively complex process. Storage of LOX for
ozone generation is subject to building and fire codes. Ozone is a
highly toxic gas and the ozone production and application facilities
must be monitored for ambient ozone. Hydrogen peroxide is a
hazardous material requiring secondary containment for storage
facilities.
7.9 References
1. Aieta, E.M., K.M. Reagan, J.S. Lang, L. McReynolds, J-W Kang, and W.H. Glaze. 1988.
"Advanced Oxidation Processes for Treating Groundwater Contaminated with TCE and PCE:
Pilot-Scale Evaluations." J. AWWA. 88(5): 64-72.
April 1999
7-15
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7. PEROXONE (OZONE/HYDROGEN PEROXIDE)
2. Bancroft, K.P., et al. 1984. "Ozonation and Oxidation Competition Values." Water Res. 18:473.
3. Chutny, B. and J. Kucera. 1974. "High Energy Radiation-induced Synthesis of Organic
Compounds. I. Introduction Isomerization and Carbon-Skeleton Changes, Radiation Synthesis
in Aqueous Solutions." Rod. Res. Rev. 5:1-54.
4. Dahi, E. 1976. "Physicochemical Aspects of Disinfection of Water by Means of Ultrasound and
Ozone." Water Res. 10:677.
5. Duguet, J., E. Brodard, B. Dussert, and J. Malevialle. 1985. "Improvement in the Effectiveness
of Ozonation of Drinking Water Through the Use of Hydrogen Peroxide." Ozone Sci. Engrg.
7(3):241-258.
6. Dukes, E.K., and M.L. Hydier. 1964. "Determination Of Peroxide By Automated Chemistry."
Anal. Chem. 36:1689-1690.
7. Ferguson, D.W., J.T. Gramith, and M.J. McGuire. 1991. "Applying Ozone for Organics Control
and Disinfection: A Utility Perspective." /. A WWA. 83(5):32-39.
8. Ferguson, D.W., M.J. McGuire, B. Koch, R.L. Wolfe, and E.M. Aieta. 1990. "Comparing
Peroxone an Ozone for Controlling Taste and Odor Compounds, Disinfection Byproducts, and
Microorganisms." J. AWWA. 82(4): 181.
9. Glaze, W. H., et al. 1987. "The Chemistry of Water Treatment Processes Involving Ozone,
Hydrogen Peroxide, and Ultraviolet Radiation." Ozone: Sci. Engrg. 9(4):335.
10. Glaze, W.H. 1986. "Reaction Products of Ozone: A Review." Environ. Health Perspectives,
69:151-157.
11. Glaze, W.H., and J-W Kang. 1988. "Advanced Oxidation Processes for Treating Groundwater
Contaminated With TCE and PCE: Laboratory Studies." J. AWWA. 88(5): 57-63.
12. Gordon, G., Cooper, W.J., Rice, R.G., and Pacey, G.E. 1992. Disinfectant Residual Measurement
Methods. Second Edition. AWWARF and AWWA, Denver, CO.
13. Hamelin, C. and Y.S- Chung. 1974. "Optimal Conditions for Mutagenesis by Ozone in
Escherichia coli Kl2." Mutation Res. 24:271.
14. Hoigne1, J. and H. Bader. 1975. "Ozonation of Water: Role of Hydroxyl Radicals as Oxidizing
Intermediates." Science. 190(4216):782.
15. Hoigne" J. and H. Bader. 1978. "Ozone Initiated Oxidations of Solutes in Wastewater: A Reaction
Kinetic Approach." Progress Water Technol. 10(516):657.
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
16. Hoigne J. and Bader. 1976. "The Role of Hydroxyl Radical Reacting in Ozonation Processes in
Aqueous Solutions." Water Resources. 10:377.
17. Huck, P..M., Anderson, W.B. Lang, C.L. Anderson, W.A. Fraser, J.C. Jasim, S.Y. Andrews, S.A.
and Pereira, G. 1995. Ozone vs; Peroxone for Geosmin and 2-Methylisoborneol Control:
Laboratory, Pilot and Modeling Studies. Conference proceedings, AWWA Annual Conference,
Anaheim, CA.
18. Ishizaki, K. et al. 1987. "Effect of Ozone on Plasmid DNA of Esherichia coli In Situ." Water
Res. 21(7):823.
19. Karimi, A.A., J.A. Redman, W.H. Glaze, and G.F. Stolarik. 1997. "Evaluation an AOP for TCE
and PCE Removal." J. AWWA. 89(8):41-53.
20. Krasner, S.W., W.H. Glaze, U.S. Weinberg, P.A. Daniel, and I.N. Najm. 1993. "Formation and
Control of Bromate During Ozonation of Waters Containing Bromide." /. AWWA. 85(1):73-81.
21. Lund, E. 1963. "Significance of Oxidation in Chemical Inactivation of Poliovirus." Arch.
Gesamite Virusforsch. 12:648.
22. Masschelein, W., M. Denis, and R. Ledent. 1977. "Spectrophotometric Determination Of
Residual Hydrogen Peroxide." Water Sewage Works. 69-72.
23. Masten, S.J. and J. Hoigne. 1992. "Comparison of Ozone and Hydroxyl Radical-Induced
Oxidation of Chlorinated Hydrocarbons in Water." Ozone Sci. Engrg. 14(3): 197-214.
24. Mentel, R. and J. Schmidt. 1973. "Investigations of Rhinoviruse Inactivation by Hydrogen
Peroxide." Acta Virol. 17:351 .
25. MWDSC and JMM (Metropolitan Water District of Southern California and James M.
Montgomery Consulting Engineers). 1992. "Pilot Scale Evaluation of Ozone and Peroxone."
AWWARF and AWWA, Denver, CO.
26. Ohlrogge, J.B. and T.P. Kernan. 1983. "Toxicity of Activated Oxygen: Lack of Dependence on
Membrane Fatty Acid Composition." Biochemical and Biophysical Research Communications.
27. Pereira, G., P.M. Huck, and W.A. Anderson. 1996. "A Simplified Kinetic Model for Predicting
Peroxone Performance for Geosmin Removal in Full-Scale Processes." Conference proceedings,
AWWA Water Quality Technology Conference; Part I. New Orleans, LA.
28. Pryor, W.A. M.M. Dooley, and D.F. Church. 1983. "Mechanisms for the Reaction of Ozone with
Biological Molecules: The Source of the Toxic Effects of Ozone." Advances in Modern
April 1999 7-17 EPA Guidance Manual
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7. PEROXONE (OZONE / HYDROGEN PEROXIDE)
Environmental Toxicology. M.G. Mustafa and M.A. Mehlman (editors). Ann Arbor Science
Publishers, Ann Arbor, MI.
29. Riesser, V. N., et al. 1977. "Possible Mechanisms for Poliovirus Inactivation by Ozone." Forum
on Ozone Disinfection. E.G. Fochtman, et al. (editors). International Ozone Institute Cleveland,
OH.
30. Scott, D.B.N. and E.G. Lesher. 1963. "Effect of Ozone on Survival and Permeability of
Escherichia coli." J. Bacteriology. 85:567.
31. Scott, K.N., et al. 1992. "Pilot-Plant-Scale Ozone and Peroxone Disinfection ofGiardia muris
Seeded Into Surface Water Supplies." Ozone Set. Engrg. 14(1):71.
32. Standard Methods. 1995. Standard Methods for the Examination of Water and Wastewater,
nineteenth edition, Franson, M.H., Eaton, A.D., Clesceri, L.S., and Greenberg, A.E., (editors).
American Public Health Association, AWWA, and Water Environment Federation, Washington
D.C.
33. Stolarik, G., and J.D. Christie. 1997. "A Decade of Ozonation in Los Angeles." Conference
proceedings, IOA Pan American Group Annual Conference, Lake Tahoe, NV.
34. Tobiason, I.E., J.K. Edzwald, O.D. Schneider, M.B. Fox, and H.J. Dunn. 1992. "Pilot Study of
the Effects of Ozone and Peroxone on In-Line Direct Filtration." /. AMWA. 84(12):72-84.
35. Von Gunten, U., A. Bruchet, and E. Costentin. 1996. "Bromate Formation in Advanced
Oxidation Processes." J. AWWA. 88(6):53.
36. Wallis, P.M. et al. 1990. "Inactivation ofGiardia Cysts in Pilot Plant Using Chlorine Dioxide and
Ozone." Conference proceedings, AWWA Water Quality Technology Conference, Philadelphia,
PA.
37. Wickramanayake, G.B. 1984. Kinetics and Mechanism of Ozone Inactivation of Protozoan
Cysts. Ph.D dissertation, Ohio State University.
38. Wolfe, R.L. et al. 1989a. "Disinfection of Model Indicator Organisms in a Drinking Water Pilot
Plant by Using Peroxone." Appl. Environ. Microbiol. 55:2230.
39. Wolfe, R.L., et al. 1989b. "Inactivation ofGiardia muris and Indicator Organisms Seeded in
Surface Water Supplies by Peroxone and Ozone." Environ. Sci. Technol. 23(6):774.
40. Yoshe-Purer, Y. and E. Eylan. 1968. "Disinfection of Water by Hydrogen Peroxide." Health
Lab Sci. 5.
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Unlike most disinfectants, ultraviolet (UV) radiation does not inactivate microorganisms by chemical
interaction. UV radiation inactivates organisms by absorption of the light which causes a
photochemical reaction that alters molecular components essential to cell function. As UV rays
penetrate the cell wall of the microorganism, the energy reacts with nucleic acids and other vital cell
components, resulting in injury or death of the exposed cells. There is ample evidence to conclude
that if sufficient dosages of UV energy reach the organisms, UV can disinfect water to whatever
degree is required. However, there has been some public health concerns with respect to the overall
efficiency of UV to disinfect potable water.
Based on the available research literature, it appears that although exceptional for disinfection of
small microorganisms such as bacteria and viruses, UV doses required to inactivate larger protozoa
such as Giardia and Cryptosporidium are several times higher than for bacteria and virus inactivation
(White, 1992; DeMers and Renner, 1992). As a result, UV is often considered in concert with ozone
and/or hydrogen peroxide to enhance the disinfection effectiveness of UV or for ground water where
Giardia and Cryptosporidium are not expected to occur.
8.1 UV Chemistry (Photochemical)
8.1.1 UV Radiation
UV radiation quickly dissipates into water to be absorbed or reflected off material within the water.
As a result, no residual is produced. This process is attractive from a DBF formation standpoint;
however, a secondary chemical disinfectant is required to maintain a residual throughout the
distribution system, which may be subject to recontamination.
UV radiation energy waves are the range of electromagnetic waves 100 to 400 nm long (between the
X-ray and visible light spectrums). The division of UV radiation may be classified as Vacuum UV
(100-200 nm), UV-C (200-280 nm), UV-B (280-315 nm) and UV-A (315-400 nm). In terms of
germicidal effects, the optimum UV range is between 245 and 285 nm. UV disinfection utilizes
either: low-pressure lamps that emit maximum energy output at a wavelength of 253.7 nm; medium-
pressure lamps that emit energy at wavelengths from 180 to 1370 nm; or lamps that emit at other
wavelengths in a high intensity "pulsed" manner.
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8.1.2 UV Disinfection Reactions
The degree to which the destruction or inactivation of microorganisms occurs by UV radiation is
directly related to the UV dose. The UV dosage is calculated as:
Where: D = UV Dose, mW-s/cm2
/ = Intensity, mW/cm2
/ = Exposure time, s
Research indicates that when microorganisms are exposed to UV radiation, a constant fraction of the
living population is inactivated during each progressive increment in time. This dose-response
relationship for germicidal effect indicates that high intensity UV energy over a short period of time
would provide the same kill as a lower intensity UV energy at a proportionally longer period of time.
The UV dose required for effective inactivation is determined by site-specific data relating to the
water quality and log removal required. Based on first order kinetics, the survival of microorganisms
can be calculated as a function of dose and contact time (White, 1992; USEPA, 1996). For high
removals, the remaining concentration of organisms appears to be solely related to the dose and
water quality, and not dependent on the initial microorganism density. Tchobanoglous (1997)
suggested the following relationship between coliform survival and UV dose:
Where: N = Effluent coliform density, /lOOmL
D = UV dose, mW-s/cm2
n = Empirical coefficient related to dose
/ = Empirical water quality factor
The empirical water quality factor reflects the presence of particles, color, etc. in the water. For
water treatment, the water quality factor is expected to be a function of turbidity and transmittance
(or absorbance).
8.1.3 Process Variables
Since UV radiation is energy in the form of electromagnetic waves, its effectiveness is not limited by
chemical water quality parameters. For instance, it appears that pH, temperature, alkalinity, and total
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inorganic carbon do not impact the overall effectiveness of UV disinfection (AWWA and ASCE,
1990). However, hardness may cause problems with keeping the lamp sleeves clean and functional.
The presence, or addition, of oxidants (e.g., ozone and/or hydrogen peroxide) enhances UV radiation
effectiveness. The presence of some dissolved or suspended matter may shield microorganisms from
the UV radiation. For instance, iron, sulfites, nitrites and phenols all absorb UV light (DeMers and
Renner, 1992). Accordingly, the absorbance coefficient is an indication of this demand and is unique
for all waters. As a result, specific "design" parameters vary for individual waters and should be
determined empirically for each application.
UV demand of water is measured by a spectrophotometer set at a wavelength of 254 nm using a 1 cm
thick layer of water. The resulting measurement represents the absorption of energy per unit depth, or
absorbance. Percent transmittance is a parameter commonly used to determine the suitability of UV
radiation for disinfection. The percent transmittance is determined from the absorbance (A) by the
equation:
Percent Transmittance = 100 x 10"A
Table 8-1 shows corresponding absorbance measurements and percent transmittance measurements
for various water qualities.
Table 8-1. Water Quality and Associated UV Measurements
Source Water Quality
Excellent
Good
Fair
Absorbance
(absorbance units/cm)
0.022
0.071
0.125
Percent
Transmittance
95%
85
75
Source: DeMers and Renner, 1992.
Continuous wave UV radiation at doses and wavelengths typically employed in drinking water
applications, does not significantly change the chemistry of water nor does it significantly interact
with any of the chemicals within the water (USEPA, 1996). Therefore, no natural physiochemical
features of the water are changed and no chemical agents are introduced into the water. In addition,
UV radiation does not produce a residual. As a result, formation of THM or other DBPs with UV
disinfection is minimal (See Section 8.5, Disinfection Byproducts of UV Radiation).
8.2 Generation
Producing UV radiation requires electricity to power UV lamps. The lamps typically used in UV
disinfection consist of a quartz tube filled with an inert gas, such as argon, and small quantities of
mercury. Ballasts control the power to the UV lamps.
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8.2.1 UV Lamps
UV lamps operate in much the same way as fluorescent lamps. UV radiation is emitted from electron
flow through ionized mercury vapor to produce UV energy in most units. The difference between'the
two lamps is that the fluorescent lamp bulb is coated with phosphorous, which converts the UV
radiation to visible light. The UV lamp is not coated, so it transmits the UV radiation generated by
the arc (White, 1992).
Both low-pressure and medium-pressure lamps are available for disinfection applications. Low-
pressure lamps emit their maximum energy output at a wavelength of 253.7 nm, while medium
pressure lamps emit energy with wavelengths ranging from 180 to 1370 nm. The intensity of
medium-pressure lamps is much greater than low-pressure lamps. Thus, fewer medium pressure
lamps are required for an equivalent dosage. For small systems, the medium pressure system may
consist of a single lamp. Although both types of lamps work equally well for inactivation of
organisms, low-pressure UV lamps are recommended for small systems because of the reliability
associated with multiple low-pressure lamps (DeMers and Renner, 1992) as opposed to a single
medium pressure lamp, and for adequate operation during cleaning cycles.
Recommended specifications for low-pressure lamps (DeMers and Renner, 1992) include:
• L-type ozone-free quartz;
• Instant start (minimal delay on startup);
• Designed to withstand vibration and shock; and
• Standard nonproprietary lamp design.
Typically, low-pressure lamps are enclosed in a quartz sleeve to separate the water from the lamp
surface. This arrangement is required to maintain the lamp surface operating temperature near its
optimum of 40°C. Although Teflon sleeves are an alternative to quartz sleeves, quartz sleeves absorb
only 5 percent of the UV radiation, while Teflon sleeves absorb 35 percent (Combs and McGuire,
1989). Therefore, Teflon sleeves are not recommended.
8.2.2 Ballasts
Ballasts are transformers that control the power to the UV lamps. Ballasts should operate at
temperatures below 60°C to prevent premature failure. Typically, the ballasts generate enough heat to
warrant cooling fans or air conditioning (White, 1992).
Two types of transformers are commonly used with UV lamps; namely, electronic and
electromagnetic. Electronic ballasts operate at a much higher frequency than electromagnetic
ballasts, resulting in lower lamp operating temperatures, less energy use, less heat production, and
longer ballast life (DeMers and Renner, 1992).
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Typical ballast selection criteria include (DeMers and Renner, 1992):
• Underwriter's Laboratories (UL) approval;
• Compatibility with UV lamps; and
• Waterproof enclosure in remote location.
8.2.3 UV Reactor Design
Most conventional UV reactors are available in two types; namely, closed vessel and open channel.
For drinking water applications, the closed vessel is generally the preferred UV reactor for the
following reasons (USEPA, 1996):
• Smaller footprint;
• Minimized pollution from airborne material;
• Minimal personnel exposure to UV; and
• Modular design for installation simplicity.
Figure 8-1 shows a conventional closed vessel UV reactor. This reactor is capable of providing UV
dosages adequate to inactivate bacteria and viruses for flows up to 600 gallons per minute. However,
it is incapable of the higher dosages required for protozoan cysts. To increase the dosage, either the
number of UV lamps and/or the exposure time should be increased.
Additional design features for conventional UV disinfection systems include:
• UV sensors to detect any drop in UV lamp output intensity;
• Alarms and shut-down systems;
• Automatic or manual cleaning cycles; and
• Telemetry systems for remote installations.
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Illumination Lamp
Monitoring Panel
Quartz Jackets
Enclosing UV Lamps
Ultrasonic Cleaning
Transducer
UV Intensity
Measuring Cell
Source: USEPA, 1996.
Figure 8-1. Closed Vessel Ultraviolet Reactor
In addition to conventional UV systems, two other UV processes are currently being evaluated for
drinking water disinfection:
• Micro-screening/UV; and
• Pulsed UV
Both of these systems profess to provide sufficient UV dose to inactivate Giardia cysts and
Cryptasporidium oocysts. See Section 8.2.3.2, Emerging UV Reactor Designs, for further discussion
on these emerging technologies.
8.2.3.1 Hydraulic Design Considerations
The major elements that should be considered in the hydraulic design of a UV closed vessel reactor
are: dispersion, turbulence, effective volume, residence time distribution, and flow rate (USEPA,
1996).
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Dispersion
Dispersion is the characteristic of water elements to scatter spatially. The ideal UV reactor is plug
flow, where water particles are assumed to discharge from the reactor in the same sequence they
entered and each element of water passing through the reactor resides in the reactor for the same
period of time. An ideal plug flow reactor has no dispersion and is approximated by a long tank with
high length-to-width ratio in which dispersion is minimal (USEPA, 1996).
Turbulence
In addition to plug flow characteristics, the ideal UV reactor has a flow that is turbulent radially from
the direction of flow, to eliminate dead zones. This radially turbulent flow pattern promotes uniform
application of UV radiation. A negative of having a radially turbulent flow pattern is that some axial
dispersion results, thus disrupting the plug flow characteristics. Techniques such as misaligning the
inlet and outlet, and using perforated stilling plates, have been used to accommodate the
contradicting characteristics of plug flow and turbulence (USEPA, 1996).
8.2.3.2 Emerging UV Reactor Designs
Two emerging technologies in UV reactor design are discussed below. All testing of these two
systems to date has been under controlled laboratory or field conditions. Both of these technologies
require demonstration of efficacy and applicability in real-world treatment operations.
Micro-Screening/UV
This unit consists of two treatment chambers, each containing a 2 Jim nominal porosity metal screen.
Each side of the screen has three 85 watt low pressure mercury lamps, with a total of six lamps per
filter. The theoretical minimum dose is 14.6 mW-sec/cm2 at the wavelength of 254 nm. The system is
designed to capture Cryptosporidium oocysts from the water onto the first screen. The first cycle is
set to establish the UV dose. The flow within the first screen chamber is then reversed to backflush to
oocysts onto the second screen where the oocysts are trapped until the preset UV dose is reached.
Using valves, the pattern of flow is designed to ensure that oocysts are temporarily captured on both
filters so that they are exposed to a total preset UV dose, which is totally independent of treated water
flowrate (Clancy et al., 1997). Johnson (1997) stated that such a system is capable of achieving total
UV doses of 8,000 mW-sec/cm2, sufficient to inactivate Giardia cysts and Cryptosporidium oocysts.
One drawback of this type of reactor is the headloss created (up to 65 feet) by the 2 jam openings in
the screen.
Pulsed UV
In the pulsed UV reactor, capacitors build up and deliver electricity in pulses to xenon flash tubes in
the center of a 2 inch diameter flash chamber through which water flows. The unit is designed to
provide microsecond pulses at 1 to 30 hertz (or per second). With each pulse, the flash tubes give off
high intensity, broad band radiation, including germicidal UV radiation, which irradiates the flowing
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water with irradiences of 75 mW-sec/cm2 at 2 cm from the flash tube surface (Clancy et al., 1997).
The UV dose can be adjusted by increasing or decreasing the frequency of the pulsing.
8.3 Primary Oses and Points of Application
The primary use of UV radiation is to inactivate pathogens to regulated levels. UV radiation is a
physical disinfectant that leaves no residual. Therefore, it should be used only as a primary
disinfectant followed by a chemical secondary disinfectant to protect the distribution system against
coliform proliferation and biofilm formation. See Section 8.4, Pathogen Inactivation and
Disinfection Efficiency" for a detailed discussion of disinfection efficiency and limitations.
The most common point of application for UV radiation is the last step in the treatment process train
just prior to the distribution system and after filtration. The use of UV disinfection has no impact on
other processes at the water treatment facility.
8.4 Pathogen Inactivation and Disinfection Efficiency
As opposed to most alternative disinfectants, UV is a physical process that requires a contact time on
the order of seconds to accomplish pathogen inactivation (Sobotka, 1993). As with any disinfectant,
UV has its limitations. For example, because it is a physical rather than a chemical disinfectant, it
does not provide a residual to control pathogen proliferation and biofilm formation in the distribution
system. When using UV for primary disinfection, some form of secondary chemical disinfectant is
required to maintain water quality within the distribution system.
8.4.1 Inactivation Mechanism
UV radiation is efficient at inactivating vegetative and sporous forms of bacteria, viruses, and other
pathogenic microorganisms. Electromagnetic radiation in the wavelengths ranging from 240 to 280
nanometers (nmj effectively inactivates microorganisms by irreparably damaging their nucleic acid.
The most potent wavelength for damaging deoxyribonucleic acid (DNA) is approximately 254 nm
(Wolfe, 1990). Other UV wavelengths, such as 200 nm, have been shown to exhibit peak absorbance
in aqueous solutions of DNA (von Sonntag and Schuchmann, 1992); however, there is no practical
application for UV inactivation of microorganisms in the wavelength range from 190 to 210 nm
(USEPA, 1996).
The germicidal effects of UV light involve photochemical damage to RNA and DNA within the
microorganisms. Microorganism nucleic acids are the most important absorbers of light energy in
the wavelength of 240 to 280 nm (Jagger, 1967). DNA and RNA carry genetic information
necessary for reproduction; therefore, damage to either of these substances can effectively sterilize
the organism. Damage often results from the dimerization of pyrimidine molecules. Cystosine
(found in both DNA and RNA), thymine (found only in DNA), and uracil (found only in RNA) are
the three primary types of pyrimidine molecules. Replication of the nucleic acid becomes very
difficult once the pyrimidine molecules are bonded together due to the distortion of the DNA helical
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structure by UV radiation (Snider et al., 1991). Moreover, if replication does occur, mutant cells that
are unable to replicate will be produced (USEPA, 1996). Figure 8-2 is a schematic of the germicidal
inactivation observed with UV radiation.
Two phenomena of key importance when using UV disinfection in water treatment are the dark
repair mechanisms and the capability of certain organisms to photoreactivate following exposure to
certain light wavelengths. Under certain conditions, some organisms are capable of repairing
damaged DNA and reverting back to an active state in which reproduction is again possible.
Typically, photoreactivation occurs as a consequence of the catalyzing effects of sunlight at visible
wavelengths outside of the effective disinfecting range. The extent of reactivation varies among
organisms. Coliform indicator organisms and some bacterial pathogens such as Shigella have
exhibited the photoreactivation mechanism; however, viruses (except when they have infected a host
cell that is itself photoreactive) and other types of bacteria cannot photoreactivate (USEPA, 1980;
USEPA, 1986; Hazen and Sawyer, 1992). Because DNA damage tends to become irreversible over
time, there is a critical period during which photoreactivation can occur. To minimize the effect of
photoreactivation, UV contactors should be designed to either shield the process stream or limit the
exposure of the disinfected water to sunlight immediately following disinfection.
UV
253.7 nm
DNA double strand
Formation of double
bond in pyrimidine
molecules inhibits
replication
Source: Tchobanoglous, 1997.
Figure 8-2. Germicidal Inactivation by UV Radiation
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8.4.2 Environmental Effects
To achieve inactivation, UV should be absorbed into the microorganism. Therefore, anything that
prevents UV from reacting with the microorganism will decrease the disinfection efficiency.
Scheible and Bassell (1981) and Yip and Konasewich (1972) reported that pH had no effect on UV
disinfection. Several factors that are known to affect disinfection efficiency of UV are:
* Chemical and biological films that develop on the surface of UV lamps;
* Dissolved organics and inorganics;
• Clumping or aggregation of microorganisms;
• Turbidity;
• Color; and
• Short-circuiting in water flowing through the UV contactor.
8.4.2.1 Chemical Films and Dissolved Organics and Inorganics
Accumulation of solids onto the surface of the UV sleeves can reduce the applied UV intensity and,
consequently, disinfection efficiency. In addition to biofilms caused by organic material, buildup of
calcium, magnesium, and iron scales have been reported (DeMers and Renner, 1992). Waters
containing high concentrations of iron, hardness, hydrogen sulfide, and organics are more susceptible
to scaling or plating (i.e., the formation of a thin coat on unit surfaces), which gradually decreases the
applied UV intensity. Scaling is likely to occur if dissolved organics are present and inorganic
concentrations exceed the following limits (DeMers and Renner, 1992):
• Iron greater than 0.1 mg/L;
• Hardness greater than 140 mg/L; and
• Hydrogen sulfide greater than 0.2 mg/L.
Figure 8-3 shows the UV dose required for inactivation of MS-2 coliphage at two pilot plants. Snicer
et al. (1996) concluded that one possible explanation for higher UV dose for the same degree of
inactivation required at pilot plant 2 could be the amount of scaling caused by higher iron
concentrations experienced at this plant. Iron concentrations at pilot plant 2 were in the range of 0.45
to 0.65 mg/L, which exceed the limit shown above.
A variety of chemical substances can decrease UV transmission (Yip and Konasewich, 1972),
including humic acids, phenolic compounds, and lignin sulfonates (Snider et al., 1991), as well as
chromium, cobalt, copper, and nickel. It has been reported that coloring agents, such as Orzan S, tea,
and extract of leaves reduce intensity within a UV contactor (Huff, 1965). In addition, iron, sulfites,
nitrites, and phenols can absorb UV (DeMers and Renner, 1992).
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8. ULTRA VIOLET RADWDON
10
20
30
40
50
60
70
80
100
UVDose(mWsec/cm;)
- Pilot 1 Water
-*- Pilot 2 Water
Source: Snicer et al., 1996. ~" ~~
Figure 8-3. UV Dose Required for Inactivation of MS-2 Coliphage
8.4.2.2 Microorganism Clumping and Turbidity
Particles can affect the disinfection efficiency of UV by harboring bacteria and other pathogens,
partially protecting them from UV radiation, and scattering UV light (see Figure 8-4). Typically, the
low turbidity of ground water results in minimal impact on disinfection efficiency. However, the
higher turbidities of surface water can impact disinfection efficiency.
Similar to particles that cause turbidity, microorganism aggregation can impact disinfection
efficiency by harboring pathogens within the aggregates and shade pathogens that would otherwise
be inactivated.
8.4.2.3 Reactor Geometry and Short Circuiting
Poor geometry within the UV contactor (which creates spacing between lamps) can leave dead areas
where inadequate disinfection occurs (Hazen and Sawyer, 1992). A key consideration to improving
disinfection is to minimize the amount of dead spaces where limited UV exposure can occur. Plug
flow conditions should be maintained in the contactor; however, some turbulence should be created
between the lamps to provide radial mixing of flow. In this manner, flow can be uniformly
distributed through the varying regions of UV intensity, allowing exposure to the full range of
available UV radiation (Hazen and Sawyer, 1992).
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As mentioned earlier, UV systems typically provide contact times on the order of seconds.
Therefore, it is extremely important that the system configuration limit the extent of short circuiting.
Particle
shading
UV light
scatter
Complete
penetration
Incomplete
penetration
Region of limited
cellular damage
Particle
shading
UV light
scatter
Complete
penetration
Incomplete
penetration
Region of limited
cellular damage
Source: Tchobanoglous, 1997.
Figure 8-4. Particle Interactions that Impact UV Effectiveness
8.4.3 Disinfection Efficacy
UV disinfection has been determined to be adequate for inactivating bacteria and viruses. Most
bacteria and viruses require relatively low UV dosages for inactivation, typically in the range of 2 to
6 mW-s/cm2 for 1-log inactivation. Protozoan [oo]cysts, in particular Giardia and Cryptosporidium,
are considerably more resistant to UV inactivation than other microorganisms. Results of several
studies investigating the ability of UV to inactivate bacteria, viruses, and protozoa, are described in
the following sections.
8.4.3.1 Bacteria and Virus Inactivation
UV doses required for bacteria and virus inactivation are relatively low. One study determined that
UV was comparable to chlorination for inactivation of heterotrophic plate count bacteria following
treatment using granular activated carbon (Kruithof et al., 1989).
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A study of the ability of UV and free chlorine to disinfect a virus-containing ground water showed
that UV is a more potent virucide than free chlorine, even after the chlorine residual was increased to
1.25 mg/L at a contact time of 18 minutes (Slade et al., 1986). The UV dose used in this study was
25 mW-s/cm2.
Table 8-2 shows results from a more recent pilot plant study (Snicer et al., 1996). As indicated by
the different UV doses to obtain the same level of inactivation, water characteristics dramatically
impact disinfection efficiency. They believed that the higher concentration of iron in pilot plant 2
water (Figure 8-3) interfered with UV or influenced the aggregation of MS-2 viral particles.
Snicer et al. (1996) also compared the susceptibility of MS-2 coliphage to hepatitis A virus, poliovirus,
and rotavirus for 10 ground water sources. Results indicated that MS-2 was found to be approximately
2 to 3 times more resistant to UV disinfection than the three human pathogenic viruses.
Table 8-2. Doses Required for MS-2 Inactivation
Log MS-2 Inactivation
1
2
3
4
5
6
Pilot Plant 1
(mW.s/cm2)
3.9
25.3
46.7
68
89.5
111.0
Pilot Plant 2
(mW.s/cm2)
15.3
39.3
63.3
87.4
111.4
135.5
Source: Snicer etal., 1996.
8.4.3.2 Protozoa Inactivation
Even though protozoa were once considered resistant to UV radiation, recent studies have shown that
ultraviolet light is capable of inactivating protozoan parasites. However, results indicate that these
organisms require a much higher dose than that needed to inactivate other pathogens. Less than 80
percent of Giardia lamblia cysts were inactivated at UV dosages of 63 mW-s/cm2 (Rice and Hoff,
1981). A 1-log inactivation of Giardia muris cysts was obtained when the UV dose was increased to
82 mW-s/cm2 (Carlson et al., 1982).
To achieve 2-log inactivation of Giardia muris cysts, a minimum ultraviolet light dose of above 121
mW-s/cm2 is needed. Karanis et al. (1992) examined the disinfection capabilities of ultraviolet light
against Giardia lamblia cysts extracted from both animals and humans (Karanis et al., 1992). Both
groups suffered a 2-log reduction at UV doses of 180 mW-s/cm2. Two important factors to consider
when determining dose requirements for Giardia inactivation are the parasite source and the growth
stage of the microorganism (Karanis et al., 1992). Figure 8-5 shows that the source of the parasites is
important in determining dose requirements. Figure 8-6 is from a study conducted in 1992 on
Acanthamoeba rhysodes inactivation (Karanis et al., 1992). These data show that the age of the
protozoa can dramatically affect the dose required to achieve a desired level of inactivation.
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80 100
UV Dose (mWa/cm2)
180
-Cysls obtained from gerbils ~B~Cysls obtained from humans |
Source: Karanis et al., 1992.
Figure 8-5. UV Doses Required to Achieve Inactivation of Giardia lamblia Cysts
Obtained from Two Different Sources
•0.25
-0.5
•075
•1,25
•1.75
10
20
30 40 50
UV Dose (mWs/cm2>
*~7-day-otd cysts ~&~28-day-old cysts |
Source: Karanis et al., 1992.
Figure 8-6. Impact of Growth Stage of A. rhysodes on the Required UV Dosage to
Achieve Inactivation
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Results from recent studies show a potential for inactivating Cryptosporidium parvum oocysts using
ultraviolet light disinfection. A 2 to 3-log reduction in the viability of Cryptosporidium parvum
oocysts was achieved using a low-pressure ultraviolet light system with a theoretical minimum
intensity of 14.58 mW/cm2 and a contact time of 10 minutes (ultraviolet dose of 8,748 mW-s/cm2)
(Campbell et al., 1995). The combination filter-UV system described by Johnson (1997) is capable
of delivering doses as high as 8,000 mW-s/cm2, sufficient to achieve 2-log Cryptosporidium oocyst
inactivation.
A pulsed UV process that delivered a minimum dose of 1,900 mW-s/cm2 to any particle within the
reactor was found to achieve Cryptosporidium oocyst inactivation levels in the range of 2-log.
(Clancy et al., 1997). In this study, the reactor residence time was 4.7 seconds and the unit was
operated to deliver 46.5 pulses per volume (10 Hz). Each pulse transfers power at the intensity rate
of about 41 mW-s/cm2.
8.4.3.3 UV Dose Summary
UV radiation is considered effective for inactivating bacterial and viral pathogens. UV doses for 2
and 3-log inaetivation of viruses are 21 and 36 mW-s/cm2, respectively (AWWA, 1991). These
doses are based on studies of hepatitis A virus inactivation and were derived by applying a safety
factor of 3 to the inactivation data. Results from a more recent study on several ground water sources
of hepatitis A virus, indicate that a 6 to 15 mW-s/cm2 dose is required for a 4-log inactivation (Snicer
et al., 1996). A safety factor was not applied to the doses. A 4-log inactivation of bacteriophage
MS-2 is achieved at a dose of 93 mW-s/cm2 (Snicer et al., 1996). In ground water containing high
iron levels (0.65 ppm), applying a safety factor of 1.5 to the highest reported dose against viruses will
yield a UV dose of 140 mWs/cm2for a minimum 4-log inactivation of viruses
In summary, ultraviolet radiation is effective against bacteria and viruses at low dosages. However,
much higher dosages are required for Cryptosporidium and Giardia inactivation.
8.5 Disinfection Byproducts of UV Radiation
Unlike other disinfectants, UV does not inactivate microorganisms by chemical reaction. However,
UV radiation causes a photochemical reaction in the organism RNA and DNA. The literature
suggests that UV radiation of water can result in the formation of ozone or radical oxidants (Ellis and
Wells, 1941; Murov, 1973). Because of this reaction, there is interest in determining whether UV
forms similar byproducts to those formed by ozonation or advanced oxidation processes.
8.5.1 Ground Water
Malley et al. (1995) analyzed 20 ground water samples for aldehydes and ketones before and after
UV radiation. Only one ground water sample, which contained 24 mg/L non-purgeable DOC and
was highly colored, contained DBFs after exposure to UV. Low levels of formaldehyde were
measure in duplicated experiments for this UV treated ground water sample mentioned above. GC-
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ECD chromatographs before and after UV radiation for the other 19 ground waters studied showed
significant shifts or unknown peaks after exposure to UV.
Malley et al. (1995) also determined the influence of UV on DBF formation during subsequent
chlorination. To examine these effects, the 20 ground water samples were subjected to simulated
distribution system (SE>S) DBF tests with chlorine, before and after UV radiation. The data indicate
that UV radiation did not significantly alter the SDS/DBP formation by chlorine in the ground waters
studied.
To examine the effects of varying UV dosages on DBF formation, six new ground water samples
(Malley et al., 1995) were subjected to UV dosages of 60, 130, and 200 mW-s/cm2. In this case,
DBFs were not formed by UV radiation for any of the ground waters tested at any of the UV
dosages. A comparison of chromatographs for samples before and after UV radiation, and for each
UV dosage, showed no significant differences or appearances of unknown peaks.
8.5.2 Surface Water
UV radiation was found to produce low levels of formaldehyde in the majority of surface waters
studied (Malley et al., 1995). The highest formaldehyde concentrations, ranging up to 14 \igfL, were
observed in UV treatment of raw water, whereas trace levels (1 to 2 (Jg/L) were found in UV
treatment of conventionally treated water. Since formaldehyde formation was also observed for one
of the ground water samples, it appears that UV radiation of waters containing humic matter (i.e.,
color producing, UV absorbing organic macromolecules) will result in low levels of formaldehyde
formation. Chromatographic examination of the surface water samples before and after UV radiation
• :M!'! " " • . ' . '; .
showed no other significant changes in the GC-ECD chromatograms.
Because of the chlorine demands of surface waters, higher chlorine dosages were required for post
disinfection following UV radiation. This resulted in larger DBF concentrations than in the ground
waters studied (Malley et al., 1995). However, the overall effect of UV radiation on SDS/DBPs was
insignificant. As in the ground water studies, UV radiation did not significantly alter the total
concentration or the speciation of the disinfection byproducts (e.g., THMs, HAAS, HANs, or HKs).
8.5.3 DBF Formation with Chlorination and Chloramination
following UV Radiation
Research results suggest that UV radiation does not directly form DBFs or alter the concentration or
species of DBFs formed by post-disinfection (Malley et al., 1995). However, the question of whether
UV radiation influences the rate of DBF formation by post-disinfection is important. Several studies
have addressed this question. Two surface waters that produced significant concentrations of a wide
variety of DBFs in previous tests were chosen as samples. With the chlorine residuals carefully
monitored to ensure they were consistent for pre-UV and post-UV samples, the results of the
experiments suggested that UV radiation did not significantly affect the rate of DBF formation.
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Studies were only performed to determine the pentane extractable DBF formation rate of a surface
water sample for varying pH conditions. The results showed that UV radiation did not affect the rate
of chloroform formation at pH 8.0 (Malley et al., 1995). Similarly, UV did not affect the DBF
formation rate at pH 5.0. At pH 8.0, chloroform was the only pentane extractable detected, whereas
atpH 5.0, chloroform, bromodichloromethane, chlorodibromomethane, and 1,1,1-trichloroacetone
were formed.
The effects of UV radiation on the DBF formation rate following chloramination were also tested in
this study using a surface water sample (Malley et al., 1995). Chloroform, trichloroacetic acid,
dichloroacetonitrile (at low levels), and cyanogen chloride (at low levels) were the only detectable
DBFs. Chloroform was the only compound formed at pH 8.0, and its rate of formation was not
affected by UV radiation. At pH 5.0, chloroform and dichloroacetonitrile were formed, but their rate
of formation was unaffected by UV radiation. Data showed that the effects of UV radiation on
cyanogen chloride formation at pH 8.0 and pH 5.0 had no significant trends.
In summary, the DBF formation rate studies indicated that UV radiation did not significantly affect
DBF formation rates when chlorine or chloramines were used as the post-disinfectant.
8.6 Status of Analytical Methods
Ultraviolet radiation intensity meter readings are taken to monitor the output of the UV disinfection
system. These readings, coupled with the flow through the UV reactor, are used to determine the UV
dosage applied. UV radiation leaves no residual disinfectant behind to monitor. Therefore, some
form of secondary chemical disinfectant should be added to protect the distribution system against
coliform proliferation and biofilm formation. Analytical methods for these chemical disinfectants are
discussed elsewhere in this report.
8.6.1 Monitoring of Generated Ultraviolet Radiation
Ultraviolet intensity at 253.7 nm (the predominant wavelength emitted by low pressure mercury
vapor lamps) is the water quality parameter used to monitor UV disinfection system output (Snider et
al., 1991). The rate of disinfection is directly related to the average intensity of the UV light. Since
UV intensity probes can only indicate UV intensity at a single point, there is no practical way to
measure the average intensity of a UV system in the field by the operator.
The average intensity is dependent upon the three dimensional lamp geometry. Scheible (1983)
developed a mathematical model that calculates the intensity at any point within the UV reactor. This
Point Source Summation (PSS) method is used to estimate the average intensity emitted by any
specific unit. UV disinfection system manufacturers use this method to design the system. The
single point UV intensity probe reading is typically used for routine monitoring of the UV system
only.
UV intensity sensors are typically photodiode sensors properly filtered to monitor the lamp intensity
in the germicidal range only (DeMers and Renner, 1992). A minimum of two sensors (mounted near
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separate lamps) located at the center of each lamp for each reactor is recommended as part of the
"system controls and instrumentation. White (1992) recommends installing the sensors in the wall of
the disinfection chamber at the point of greatest distance from the tube or tubes.
The UV sensors should continuously sense the UV intensity produced in the bank of lamps. The
sensor should be field calibrated to account for lamp geometry. Each UV sensor installed should
provide a "low UV output" warning and a "low-low UV output" alarm indication that is field
adjustable.
8.6.1.1 Ultraviolet Sensor Performance
To test electronic sensor performance, Snicer et al. (1996) placed a single electronic UV sensor at the
center of a UV reactor. The UV sensor converted the UV energy into an electronic signal, which was
used to indicate system performance. Initial results indicated that the sensor originally installed with
the two UV reactors tended to wear or degrade in performance over time. The original sensor
readings in the pilot facility were erratic and had a general downward trend over 6 months of
operation. This loss in performance could not be attributed to UV lamp aging. In addition, these
readings did not correlate well to actual system performance. After operation of these original
sensors for 6 months, the manufacturer was consulted and a new type of sensor was installed into
both pilot facilities. The new sensor design specifically addressed the problems encountered during
the initial 6 months of the study. These new proprietary sensors performed! consistently and sensor
readings were shown to correlate well with actual MS-2 coliphage inactivation.
8.6.2 Disinfectant Interferences
Suspended solids may be the most important water quality parameter impacting UV intensity
measurements. Particles can harbor bacteria and at least partially protect them from UV light.
Particles can be completely penetrated, partially penetrated, or scatter UV light (Figure 8-4). All
particles in water may not absorb UV light. Quails et al. (1983) suggested that clays merely scatter
UV light and, therefore, do little to inhibit performance.
Yip and Konasewich (1972) listed many chemical substances that interfere with UV transmission at
253.7 nm including phenolic compounds, humic acids, and ferric iron.
8.7 Operational Considerations
Onsite pilot plant testing is recommended to determine the efficiency and adequacy of UV
disinfection for a specific quality of water. The efficiency test involves injecting select
microorganisms into influent water and sampling effluent water to determine survival rates. The
National Science Foundation's Standard 55 for ultraviolet water treatment systems recommends that
UV disinfection systems not be used if the UV transmittance is less than 75 percent (NSF, 1991). If
the raw water UV transmittance is less that 75 percent, the UV system should be proceeded by other
treatment processes (to increase UV transmittance) or a different disinfectant should be used.
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As previously discussed, some constituents that adversely interfere with UV disinfection
performance by either scattering and/or absorbing radiation are iron, chromium, copper, cobalt,
sulfites, and nitrites. Care should be taken with chemical processes upstream of UV disinfection
process to minimize increasing concentrations of these constituents since disinfection efficiency may
be adversely affected.
8.7.1 Equipment Operation
UV disinfection facilities should be designed to provide flexibility in handling varying flow rates.
For lower flow rates, a single reactor vessel should be capable of handling the entire flow rate. A
second reactor vessel with equal capacity of the first reactor vessel should be provided for
redundancy should the first reactor vessel be taken out of service. For higher flow rates, multiple
reactor vessels should be provided with lead/lag operation and flow split capacity to balance run time
for each reactor vessel, if desired, and to avoid hydraulic overloading. Valves should be provided
within the interconnecting piping to isolate one reactor vessel from another. There should also be a
positive drainage system to remove water from within a reactor vessel when it is taken out of service.
8.7.1.1 UV Lamp Aging
The output of UV lamps diminishes with time. Two factors that affect their performance are:
solarization which is the effect UV radiation has on the UV lamp that causes it to become opaque;
and, electrode failure which occurs when electrodes deteriorate progressively each time the UV lamp
is cycled on and off. Frequent lamp cycling will lead to premature lamp aging. When determining
the requirement for UV disinfection, a 30 percent reduction of UV output should be used to estimate
end of lamp. Average life expectancy for low pressure UV lamps is approximately 8,800 hours.
8.7.1.2 Quartz Sleeve Fouling
Fouling of the quartz sleeve reduces the amount of UV radiation reaching the water. The quartz
sleeve has a transmissibility of over 90 percent when new and clean. Over time, the surface of the
quartz sleeve that is in contact with the water starts collecting organic and inorganic debris (e.g., iron,
calcium, silt) causing a reduction in transmissibility (USEPA, 1996). When determining the
requirements for UV disinfection, a 30 percent reduction of UV transmission should be used to
reflect the effect of quartz sleeve fouling.
8.7.2 Equipment Maintenance
8.7.2.1 UV Lamp Replacement
Adequate space should be provided around the perimeter of the reactor vessels to allow access for
maintenance and replacement of UV lamps. With modular electrical fittings, lamp replacement
consists of unplugging the pronged connection of the old lamp and plugging in the new.
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8.7.2.2 Quartz Sleeve Cleaning
Quartz sleeve cleaning may be accomplished by physical or chemical means. Physical alternatives
include:
• Automatic mechanical wiper;
• Ultrasonic devices;
* High water pressure wash; and
» Air scour.
Chemical cleaning agents include sulfuric or hydrochloric acid. A UV reactor vessel may contain
one of more physical cleaning system with provision for an occasional chemical cleaning.
8.7.2.3 Miscellaneous
I ' "', * '• !,'| '' , '
Effective maintenance of a UV system will involve:
• Periodic checks for proper operation;
• Calibration of intensity meter for proper sensitivity; and
• Inspect and/or clean reactor vessel interior.
8.7.3 Standby Power
Producing UV radiation requires electricity to power the electronic ballasts, which in turn power the
UV lamps. Since disinfection is of utmost importance in producing potable water, the UV system
should remain in service during periods of primary power failure. A dual power feed system or
essential circuitry powered by a standby generator are typical ways to achieve the desired reliability.
Each low pressure UV lamp requires approximately 100 Watts of standby power. A second
precaution that should be considered is not powering the UV system from the same motor control
center (MCC) that powers variable frequency drives (VFDs). The electronic ballasts produce
harmonics that may require mitigation (active harmonic filters) for the VFDs.
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8.8 Summary Table
Table 8-3. Summary of UV Disinfection
Consideration
Generation
Primary uses
Inactivation efficiency
Byproduct formation
Limitations
Point of application
Special considerations
Description
Low pressure and medium pressure UV lamps available.
Primary physical disinfectant; requires secondary chemical
disinfectant for residual in distribution system.
Very effective against bacteria and viruses at low dosages
(5-25 mW»s/cm2 for 2-log removal and 90-140 mW»s/cm2for 4-log
removal). Much higher dosage required for Cryptosporidium and
Giardia (100-8,000 mW«s/cm2 for 2-log removal).
Minimal disinfection byproducts produced.
Limited experience and data with flows greater than 200 GPM.
Water with high concentrations of iron, calcium, turbidity, and
phenols may not be applicable to UV disinfection.
Prior to distribution system.
Extremely high UV dosages for Cryptosporidium and Giardia may
make surface water treatment impractical.
8.9 References
1. AWWA (American Water Works Association). 1991. Guidance Manual for Compliance with
the Filtration and Disinfection Requirements for Public Water Systems Using Surface Water
Sources.
2. AWWA and ASCE (American Society of Civil Engineers). 1990. Water Treatment Plant Design.
Second edition, McGraw-Hill, Inc., New York, NY.
3. Campbell, A.T., et al. 1995. "Inactivation of Oocysts of Cryptosporidium parvum by Ultraviolet
Radiation." Water Res. 29(11):2583.
4. Carlson, D. A., et al. 1982. Project Summary: Ultraviolet Disinfection of Water for Small Water
Supplies. Office of Research and Development, U.S. Environmental Protection Agency;
Cincinnati, OH, EPA/600/S2-85/092.
5. Clancy, J.L., T.M. Hargy, M.M. Marshall, and I.E. Dyksen. 1997. "Inactivation of
Cryptosporidium parvum Oocysts in Water Using Ultraviolet Light." Conference proceedings,
AWWA International Symposium on Cryptosporidium and Cryptosporidiosis, Newport Beach,
CA.
April 1999
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6. Combs, R., and P. McGuire. 1 989. "Back to Basics - The Use of Ultraviolet Light for Microbial
Control." Ultrapure Water Journal. 6(4):62-68.
7. E)eMers, L.D. and R.C. Renner. 1992. Alternative Disinfection Technologies For Small Drinking
Water Systems. AWWARF.
8. Ellis, C., and A. A. Wells. 1 94 1 . The Chemical Action of Ultraviolet Rays. Reinhold Publishing
Co., New York, NY.
9. Hazen and Sawyer. 1992. Disinfection Alternatives for Safe Drinking Water. Van Nostrand
Reinhold, New York, NY.
10. Huff, C. B. 1965. "Study of Ultraviolet Disinfection of Water and Factors in Treatment
Efficiency." Public Health Reports. 80(8):695-705.
1 1. Jagger, J. 1967. Introduction to Research in Ultraviolet Photobiology. Prentice-Hall Inc.,
s, NJ.
12. Johnson, R.C. 1997. "Getting the Jump on Cryptosporidium with UV." Opflow. 23(10):!.
13. Karanis, P., et al. 1992. "UV Sensitivity of Protozoan Parasites." J. Water Supply Res. Technol.
14. Kruithof, J.C., et al. 1989. Summaries, WASSER BERLIN '89; International Ozone Association,
European Committee, Paris.
15. Malley Jr., J.P, J.P. Shaw, and J.R. Ropp. 1995. " Evaluations of Byproducts by Treatment of
Groundwaters With Ultraviolet Irradiation." AWWARF and AWWA, Denver, CO.
16. Murov, S.L. 1973. Handbook of Photochemistry. Marcel Dekker, New York, NY.
17. IjSF (National Science Foundation). 1991. NSF Standard 55: Ultraviolet Water Treatment
Systems. National Sanitation Foundation, Ann Arbor, MI.
18. Quails, R., Flynn, M., and Johnson, J. 1983. "The Role of Suspended Particles in Ultraviolet
Disinfection." /. Water Pollution Control Fed. 55( 1 0): 1 280- 1 285.
19. Rice, E.W. and J.C. Hoff. 1981. "Inacti vation of Giardia lamblia Cysts by Ultraviolet
Irradiation." Appl. Environ. Microbiol. 42:546-547.
20. Scheible, O.K. and C.D. Bassell. 1981. Ultraviolet Disinfection Of A Secondary Wastewater
Treatment Plant Effluent. EPA-600/2-8 1-152, PB81-242125, U.S. Environmental Protection
Agency; Cincinnati, OH.
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21. Scheible, O.K. 1983. "Design and Operation of UV Systems." Presented at Water Pollution
Control Federation Annual Conference, Cincinnati, OH.
22. Slade, J. S., N.R. Harris, and R.G. Chisholm. 1986. "Disinfection of Chlorine Resistant
Enteroviruses in Ground Water by Ultraviolet Radiation." Water Sci. Technol. 189(10):! 15-123.
23. Snicer, G.A., J.P. Malley, A.B. Margolin, and S.P. Hogan. 1996. "Evaluation of Ultraviolet
Technology in Drinking Water Treatment." Presented at AWWA Water Quality Technology
Conference, Boston, MA.
24. Snider, K.E., J.L. Darby, and G. Tchobanoglous. 1991. Evaluation of Ultraviolet Disinfection
For Wastewater Reuse Applications In California. Department of Civil Engineering, University
of California, Davis.
25. Sobotka, J. 1993. "The Efficiency of Water Treatment and Disinfection by Means of Ultraviolet
Radiation." Water Sci. Technol. 27(3-4) :343-346.
26. Tchobanoglous, G.T. 1997. "UV Disinfection: An Update." Presented at Sacramento Municipal
Utilities District Electrotechnology Seminar Series. Sacramento, CA.
27. USEPA (U.S. Environmental Protection Agency). 1996. Ultraviolet Light Disinfection
Technology in Drinking Water Application - An Overview. EPA 81 l-R-96-002, Office of Ground
Water and Drinking Water.
28. USEPA. 1986. Design Manual: Municipal Wastewater Disinfection. EPA/625/1-86/021, Office
of Research and Development, Water Engineering Research Laboratory, Center for
Environmental Research Information, Cincinnati, OH
29. USEPA. 1980. Technologies for Upgrading Existing and Designing New Drinking Water
Treatment Facilities. EPA/625/4-89/023, Office Drinking Water.
30. Von Sonntag, C. and H. Schuchmann. 1992. "UV Disinfection of Drinking Water and By-
Product Formation - Some Basic Considerations." J. Water SRT-Aqua. 41(2):67-74.
31. White, G.C. 1992. Handbook of Chlorination and Alternative Disinfectants. Van Nostrand
Reinhold, New York, NY.
32. Wolfe, R.L. 1990. "Ultraviolet Disinfection of Potable Water." Environ. Sci. Tech. 24(6) :768-
773.
33. Yip, R.W. and D.E. Konasewich. 1972. "Ultraviolet Sterilization Of Water - Its Potential And
Limitations." Water Pollut. Control (Canada). 14:14-18.
April 1999 8-23 EPA Guidance Manual
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9. COMBINED DISINFECTANTS
Multiple disinfectants, the sequential or simultaneous use of two or more disinfectants, have been
used with increasing frequency in recent years. This trend is attributed to the fact that:
• Less reactive disinfectants, such as chloramines, have proven to be quite effective in reducing
DBFs formed during disinfection and are more effective for controlling biofilms in the
distribution system.
• Regulatory and consumer pressure to produce water that has been disinfected to achieve high
inactivation for various pathogens, has pushed the industry towards more effective disinfectants.
Sometimes more effective disinfection meant using higher disinfectant doses which also
produces more DBFs.
• Recent research has shown that the application of sequential disinfectants is more effective than
the added effect of the individual disinfectants. This process where two (or more) disinfectants
produce a synergistic effect by either simultaneous or sequential application to achieve more
effective pathogen inactivation, is referred to as interactive disinfection in this manual.
This chapter discusses recent industry applications of multiple disinfectants to meet the varied
requirements for inactivation and reduction in DBFs. The initial discussion focuses on traditional
disinfectants used in primary and secondary application. This is followed by a discussion of
interactive disinfectants where two disinfectants are applied together specifically to achieve primary
disinfection. Note that the IESWTR does not have any provision for additional credits for interactive
disinfection or taking additional credit for the synergistic effects from interactive disinfection. Until
such credit is established, interactive disinfection is considered an emerging technology. This
chapter does not discuss mixed oxidant systems, which are designed to generate mixed oxidants on-
site for drinking water disinfection, and are also considered an emerging technology.
9.1 Primary and Secondary Disinfectants
By separating the inactivation function and residual disinfection function in water treatment, each
can be optimized independently. Thus, the combination of disinfectants currently used in
disinfection is typically identified as a primary or secondary disinfectant, as follows:
• Primary disinfection refers to the inactivation of microorganisms to meet the regulatory
bacteriological requirements. This requirement typically is met by achieving certain CT
requirements to assure a target log inactivation of target organisms as set forth in the Surface
Water Treatment Rule (SWTR) (AWWA, 1991).
• Secondary disinfection refers to application of a disinfectant to meet regulatory requirements for
distribution system bacteriological quality as set forth in the Total Coliform Rule (TCR). The
SWTR requires that a residual disinfectant be measured in the distribution system, or that the
bacteriological quality meet certain standards (heterotrophic plate count (HFC) less than
500/lOOmL).
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9. COMBINED DISINFECTANTS
To be an effective primary disinfectant, the disinfectant should effectively inactivate the target
organism. Table 9-1 lists potential primary disinfectants as discussed in Chapters 2 through 8.
Table 9-1. Potential Primary Disinfectants
Potential Primary Disinfectants
Target Organism With Filtration1 Without Filtration
Coliform Bacteria Chlorine Chlorine
Chloramines Chlorine dioxide
Chlorine dioxide Interactive disinfection3
Ozone
UV
Interactive disinfection
Giardia cysts Chlorine2 Chlorine2 ~ """ "
Chlorine dioxide2 Chlorine dioxide2
Ozone2 Interactive disinfection3
Interactive disinfection
Viruses Chlorine2 Chlorine2
Chlorine dioxide2 Chlorine dioxide2
Ozone2 UV2
Pv Interactive disinfection3
Interactive disinfection
Cryptosporidium oocysts Chlorine dioxide Chlorine dioxide
Ozone Interactive disinfection3
_ Interactive disinfection
1 Natural or treatment filtration reduces disinfection inactivation requirements.
2 Inactivation credit established in SWTR.
3 Any interactive disinfection that uses ozone or peroxone without filtration is strongly discouraged.
>' • Mil!! , ' I"
As discussed in earlier chapters, certain disinfectants (e.g., ozone, UV, peroxone, and in some cases
chlorine dioxide), while being effective disinfectants, do not leave a long-lasting residual
disinfectant. Therefore, secondary disinfection is limited to those disinfectants that remain stable in
the distribution system. In order of decreasing stability, these secondary disinfectants are
chloramines, chlorine, and chlorine dioxide.
Based on the above, the combinations of disinfectants that are viable options to meet the disinfection
requirements can be determined for various treatment trains. These combinations are shown for
various treatment objectives. Note that the treatment objectives are dependent on the treatment
currently in place.
|, ' I ' , '':',:'
To meet DBF, and specifically, THM limits, several studies have evaluated the application of various
primary/secondary disinfectants. Table 9-2 presents the typical application of these combined
disinfectants.
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9. COMBINED DISINFECTANTS
Table 9-2. Primary/Secondary Disinfectant Combinations and Typical
Applications in Water Treatment
Primary / Secondary
Typical application*
Comment
Chlorine/Chlorine
Chlorine/Chloramine
Chlorine dioxide/Chlorine
dioxide
Low THMFP raw water, low TOC,
conventional treatment with
optimal coagulation.
Moderate THM production
situation, typically with
conventional treatment.
Most commonly used disinfection
scheme. Effective system.
Chlorine to provide disinfection
and monochloramine to limit DBF
formation.
High DBF production, require filter
process to remove
Cryptosporidium, low chlorine
dioxide demand in treated water.
Chlorine dioxide/Chloramine
Ozone/Chlorine
Ozone/Chloramine
UV/Chlorine
UV/Chloramine
High DBF production, require
filtration to remove
Cryptosporidium.
Moderate DBF formation, direct or
no filtration, low THMFP.
Moderate DBF formation, direct or
no filtration, higher THMFP.
Requires membrane treatment to
provide effective Giardia and
Cryptosporidium removal. UV
only for virus inactivation; ground
water disinfection; low THMFP.
Requires membrane treatment to
provide effective Giardia and
Cryptosporidium removal. UV
only for virus inactivation; ground
water disinfection, moderate
THMFP.
Primary and secondary usage
requires a limit on chlorine dioxide
dose to reduce residual
chlorate/chlorite.
Primary chlorine dioxide dose
limited to residual
chlorate/chlorite. Stable, low
reactive secondary disinfectant.
Highly effective disinfection to
achieve high log inactivation; low
THMFP to accept free chlorine.
Highly effective disinfection to
achieve high log inactivation, low
THMFP to require combined
chlorine residual.
Rare application but feasible in
special circumstances. Little
Giardia and no Cryptosporidium
inactivation.
Rare application but feasible in
special circumstances. No
Giardia or Cryptosporidium
inactivation.
Notes:
* Low DBF formation is defined as producing less than the Stage 2 D/DBP Proposed Standard (less than 0.040 mg/L TTHM; less
• than 0.030 mg/L HAAS). Moderate DBF formation is defined as producing less than the Stage 1 D/DBP Standard and more
than the Stage 2 D/DBP Proposed Standard. High DBP formation is defined as producing more than the Stage 1 D/DBP
Standard (greater than 0.080 mg/L TTHM; greater than 0.060 mg/L HAAS).
9.1.1 DBP Formation with Various Primary and Secondary
Disinfectant Combinations
The concentrations and types of DBFs formed depend on, among other things, the combination of
disinfectants used to achieve primary and secondary disinfection and the water quality. For example,
under certain water quality conditions, ozone/chloramine disinfection is known to produce lower
THM concentrations than chlorine/chloramine disinfection. However, the ozone/chloramine
alternative can increase the formation of other DBFs such as aldehydes and BOM. No single
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combination of disinfectants is applicable to all situations. Table 9-3 summarizes the potential DBFs
formed by various combinations of disinfectants. The disinfection byproducts referenced here are
discussed in greater detail in earlier chapters of this manual.
Table 9-3. DBFs Associated with Various Combined Oxidation/Disinfection
Processes
Alternative
Preoxldatlon
Chlorine
CMofJne
Chlorine dioxide
Chlorine dioxide
Potassium
permanganate
Potassium
permanganate
Ozone
Ozone
Primary
Disinfection
Chlorine
Chlorine
Chlorine
dioxide
Chlorine
dioxide
Chlorine
Chlorine
Ozone
Ozone
Secondary
Disinfection
Chlorine
Chloramine
Chlorine
Chloramine
Chlorine
Chloramine
Chlorine
Chloramine
Potential DBPs
XDBPs*
Aldehydes
XDBPs
Cyanogen chloride
Cyanogen bromide
Aldehydes
XDBPs
Aldehydes, carboxylic acids, maleic
acids
Chlorate
Chlorite
XDBPs
Aldehydes, carboxylic acids, maleic
acid
Chlorate
Chlorite
XDBPs
Aldehydes
XDBPs
Cyanogen chloride
Cyanogen bromide
Aldehydes
XDBPs
Bromate, Aldehydes, carboxylic
acids
XDBPs
Cyanogen chloride
Cyanogen bromide
Bromate, Aldehydes, carboxylic
acids
Remarks
Maximum XDBP formation compared
to all other strategies. Principal
components are TTHMs and HAAs.
Formed at relatively low levels.
Formation of XDBPs (specifically
TTHMs and HAASs) significant/y
reduced compared to chlorine/
chlorine/ chlorine.
Formed at relatively low levels.
Formation of XDBPs may be
decreased by delaying the point of
chlorine addition.
Formed at relatively low levels.
Chlorite is a major breakdown
product of chlorine dioxide.
Formation of XDBPs (especially
TTHMs and HAASs) minim/zed by
avoiding use of free chlorine.
Formed at relatively low levels.
Chlorite is a major breakdown
product of chlorine dioxide.
Formation of XDBPs may be
decreased by delaying the point of
chlorine addition.
Formed at relatively low levels.
• Formation of XDBPs may further be
decreased compared to potassium
permanganate/ chlorine/ chlorine.
Formed at relatively low levels.
Formation of certain XDBPs may
increase or decrease compared to
chlorine/ chlorine/ chlorine.
Brominated byproducts may be of
concern when bromides are present
in the raw water.
Although formed at relatively high
levels significant amounts of this
BOM can be removed through
biological filtration.
Formation of XDBPs (especially
TTHMs) minimized by avoiding use
of free chlorine.
Although formed at relatively high
levels significant amounts of this
BOM can be removed through
biological filtration.
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9. COMBINED DISINFECTANTS
, Alternative
Preoxidation
Peroxone
Peroxone
Chlorine
Potassium
Permanganate
Primary
Disinfection
Chlorine or
Ozone
Chlorine or
Ozone
UV"
UV"
Secondary
Disinfection
Chlorine
Chloramine
Chloramine
Chloramine
Potential DBPs
XDBPs
Bromate, Aldehydes, carboxylic ,
acids
XDBPs
Cyanogen chloride
Cyanogen bromide
Bromate, Aldehydes, carboxylic
acids
XDBPs
Cyanogen chloride .
Cyanogen bromide
Aldehydes
XDBPs
Aldehydes, carboxylic acids
Remarks
Formation of certain XDBPs may
increase or decrease compared to
chlorine/ chlorine/ chlorine.
Although formed at relatively high
levels significant amounts of this
BOM can be removed through
biological filtration. Also, the
formation of bromate will increase if
peroxone is used.
Formation of XDBPs may decrease
compared to peroxone/ chlorine/
chlorine.
Although formed at relatively high
levels significant amounts of this
BOM can be removed through
biological filtration. Also, the
formation of bromate will increase if
peroxone is used.
May form XDBP from pre-oxidation.
Low levels.
Very low due to less reactive
oxidants.
Very low, if any, due to less reactive
oxidants.
* XDBPs - Halogenated Disinfection Byproducts.
* * Although "conventional" UV use as primary disinfectant is limited to virus inactivation (may require membrane filtration), pulsed
UV may be able to inactivate Giardia and Cryptosporidium.
Source: Adopted in part from USEPA, 1992; Richardson etal., 1994.
Raw water chlorination, applied prior to natural organic matter (NOM) removal processes, combined
with chlorination for residual disinfection produces the greatest concentrations of halogenated DBPs.
Studies indicate that pre-oxidation of raw water with ozone or chlorine dioxide can reduce the
formation of halogenated DBPs because it shifts the point of chlorine application from raw water to
settled or filtered water which has lower DBF precursor concentrations (MWDSC and JMM, 1989).
The use of ozone can reduce the formation of halogenated byproducts in waters containing low
concentrations of bromide. However, ozone increases BOM and may encourage bacterial growth in
the distribution system. Removal of AOC with biological filtration (e.g., biological activated carbon)
reduces the potential for bacterial growth in the distribution system. The use of chloramines as a
secondary disinfectant instead of chlorine shortens the chlorine contact time and thus reduces the
formation of chlorinated byproducts. However, chloramine produces by-products of its own
(cyanogen chloride and cyanogen bromide). In addition, a short chlorine contact time prior to
ammonia addition will help inactivate heterotrophic plate count bacteria that are found in the effluent
of a biologically active filter. Bench or pilot studies will be required to evaluate the trade-offs in
DBP formation'for various disinfectant combinations for a specific application.
The application of ozone should be carefully considered because it produces aldehydes,
aldoketoacids, and carboxylic acids. However, these can be removed in a biologically active filter.
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In bromide-containing waters, ozonation can increase the formation of brominated organic DBFs and
form bromate.
In pilot plant studies for water containing low concentrations of bromide, Lykins et al. (1991)
determined that ozonation followed by chloramination produced the lowest levels of halogenated
disinfection byproducts. However, this is not applicable to source waters containing significant
bromide concentrations due to the potential for bromate formation and brominated THMs and HAAs.
The addition of chlorine dioxide will produce chlorite and chlorate and may form some oxygenated
DBFs (e.g., maleic acids).
9.1.2 Impact of Modifying Disinfection Practices
EPA and the Association of Metropolitan Water Agencies funded a 2-year study of 35 water
treatment facilities to evaluate DBF production. Among four of the facilities, eleven alternative
disinfection strategies were instigated to evaluate the difference in DBF production from the plants'
previously existing disinfection strategies. Three reports (MWDSC and JMM, 1989; Jacangelo,
1989; USEPA, 1992) analyzed and documented different aspects of the study. Table 9-4 shows the
eleven potential strategies used for primary and secondary disinfection. Table 9-5 shows the changes
in DBF production observed in the four plants after eight of these new strategies were implemented.
Following are overviews of the potential implications of these strategies, as detailed by the three
reports.
Table 9-4. Strategies for Primary and Secondary Disinfectants
Base Disinfection Condition Modified Disinfection Practice
Chlorine/Chlorine Chlorine/Chloramine
Chlorine/Chlorine Chloramine/Chloramine
Chlorine/Chlorine Chlorine Dioxide/Chloramine
Chlorine/Chlorine Ozone/Chlorine
Chlorine/Chlorine Ozone/Chloramine
Chlorine/Chlorine Chlorine Dioxide/Chlorine
Chlorine/Chlorine Chlorine Dioxide/Chlorine Dioxide
Chlorine/Chloramine Ozone/Chloramine
Chlorine/Chloramine Chlorine Dioxide/Chloramine
Ozone/Chlorine Ozone/Chloramine
Chloramine/Chloramine Ozone/Chloramine
Note: Disinfectants are listed as primary disinfectant/secondary disinfectant
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9. COMBINED DISINFECTANTS
An operational consideration of the ozone/chlorine system is the application point for chlorine. The
ozone should be completely decomposed or chemically quenched prior to chlorine addition. If ozone
is present when chlorine is added, the ozone will react with the chlorine and NOM present to form
chlorinated DBFs. ,
Ozonation converts NOM into low molecular weight humic NOM and may increase the
concentrations of precursors to some DBFs. For instance, ozonation followed by chlorination as a
secondary disinfectant may yield high concentrations of chloral hydrate (Logsdon et al., 1992;
McKnight and Reckhow, 1992). This may occur because the byproduct of ozonation, acetaldehyde,
is a known precursor for chloral hydrate, a byproduct of chlorination. Enhancement of chloral
hydrate has not been observed when biologically active filtration is used following ozonation and
' prior to chlorination (Singer, 1992). In addition to chloral hydrate, ozonation followed by
chlorination can produce greater THM and haloketones levels than chlorination alone, particularly
, when chlorine is applied at high pH levels (Jacangelo et al., 1989; Reckhow et al., 1986). Ozonation
followed by'chlorination or chloramination can increase chloropicrin and cyanogen chloride
concentrations above those observed with chlorination or chloramination alone (Jacangelo et al.,
1989). The most promising treatment strategy for preventing the enhancement of these
biodegradable ozonation byproducts and BOM is to locate ozonation after sedimentation and follow
it by biologically active GAC.
9.1.5 Chlorine/Chlorine to Ozone/Chloramine
In addition to the concerns addressed in Sections 9.1.3 and 9.1.4, switching from chlorine to
chloramine residual exposes the consumer to a residual that may be a more significant health concern
(particularly for kidney dialysis patients). The impact of switching from chlorine/chlorine to
ozone/chloramines on the production of byproducts was investigated in a 5 gpm pilot study
(MWDSC and JMM, 1989; Jacangelo et al., 1989). That switch produced greater concentrations of
chloropicrin, cyanogen chloride, formaldehyde and total aldehydes than in the original
chlorine/chlorine strategy. Concentrations of TTHMs, total haloacetic acids, total haloacetonitriles,
total haloketones and chloral hydrate were lower with ozone/chloramine. Brominated DBFs were
not reported. Ozonation followed by chloramination has been observed to increase cyanogen
chloride concentrations beyond those observed with chlorination only (Jacangelo et al., 1989).
Increased chloral hydrate has not been observed when monochloramine is applied as the secondary
disinfectant (Singer, 1992).
9.1.6 Chlorine/Chlorine to Chlorine Dioxide/Chlorine
Use of chlorine dioxide as a pre-oxidant to replace chlorine may allow moving the point of
chlorination downstream in the process train for application to water with lower NOM
concentrations. The reduced precursor concentration and the reduced chlorine dose should result in a
reduction of chlorinated DBFs. However, if excess chlorine is present in the chlorine dioxide feed
stream, it would react with NOM prior to removal in sedimentation and filtration if pre-oxidation is
practiced. Switching from chlorine/chlorine to chlorine dioxide/chlorine produces mixed results.
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Like ozone, chlorine dioxide alters the nature of NOM molecules, potentially forming greater
precursor concentrations for some DBFs while reducing the precursor concentrations for other DBFs.
The human health implications of these trade-offs are largely unknown. Chlorine dioxide/chlorine
appears to be most effective in decreasing chlorinated DBFs when it can replace the need for pre-
chlorination. However, for facilities that use pre-chlorination but do not require it, continuing to use
chlorine/chlorine while eliminating pre-chlorination may be as effective in decreasing chlorinated
DBFs.
9.1.7 Chlorine/Chlorine to Chlorine Dioxide/Chlorine Dioxicta
The potential to apply chlorine dioxide as both a primary and secondary disinfectant is limited
because:
• Chlorine dioxide is a strong oxidant and dissipates rapidly in both raw and treated waters;
and
• Approximately 50 to 70 percent of chlorine dioxide is converted to the inorganic byproducts
chlorite and chlorate.
On the positive side, chlorine dioxide/chlorine dioxide application will significantly lower the
formation of organic DBFs.
9.1.8 Chlorine/Chloramine to Ozone/Chloramine
In addition to the concerns raised in Sections 9.1.4 and 9.1.5, switching from chlorine/chloramine to
ozone/chloramine resulted in reduced formation of most of the halogenated DBFs (MWDSC and
JMM, 1989). Other studies also indicate reduction in the formation of most halogenated DBFs but
increased formation of 1,1-dichloropropanone (MWDSC and JMM, 1989). The primary difference
in chlorinated DBF formation when switching from chlorine/chloramine to ozone/chloramine could
be attributed to the shorter contact time with free chlorine.
9.1.9 Chlorine/Chloramine to Chlorine Dioxide/Chloramine
The Louisville Water Company evaluated the feasibility of switching from a chlorine/chloramine to
chlorine dioxide followed by chloramine to control THM formation (Hubbs et al., 1981). The
treatment plant includes lime soda-ash for softening. Disinfection occurs prior to the lime treatment
step. Ammonia is added to form chloramine before the water enters the softening phase. There is a
10 minute lag period between the first disinfectant (chlorine or chlorine dioxide) and second
disinfectant (chloramine) addition. The study showed a significant decrease in THM formation from
25 (0.g/L with chlorine to 5 ng/L using chlorine dioxide as the initial disinfectant. At the same time,
treated water coliform densities were essentially unchanged; however, results showed slightly more
scattered data during the chlorine dioxide test period. Based on these results, the Water Company
decided to use chlorine/chloramine to meet disinfection and THM targets. No other DBFs were
measured during the test period.
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9.; COMBINED DISINFECTANTS
9.1.10Ozone/Chlorine to Ozone/Chloramine
In addition to concerns raised in Sections 9.1.4 and 9.1.5, when compared with ozone/chlorine,
ozone/chloramine produced greater concentrations of cyanogen chloride. Concentrations of TTHM,
total haloacetic acids, total haloacetonitriles, total haloketones, chloral hydrate, total aldehydes and
formaldehyde were lower with ozone/chloramine than with ozone/chlorine. Ozone/chloramine
produces some chlorinated DBFs at greater concentrations than ozone/chlorine; however,
ozone/chloramine significantly reduces TTHMs compared to ozone/chlorine (LeLacheur et al.,
1991).
9.1.11 Summary
EPA is not encouraging systems to switch to different disinfectants due to unknown risks to public
health. When needed for compliance with regulations or increasing Cryptosporidium inactivation,
careful selection of alternative disinfectants as primary and secondary disinfectants, can produce less
DBFs and increase inactivation. In general, the results followed the characteristic DBFs associated
with the primary disinfectant (halogenated DBFs with chlorine compounds or ozone in the presence
of bromide oxidized organics, AOC with ozone or peroxone). However, by carefully selecting the
primary and secondary disinfectant and avoiding long contact times and high dosages of halogens
(chlorine, bromine), the total DBF formation declined. The quantity and types of DBFs that form are
site-specific, depending on the water quality, disinfectant dose and type, and are best determined by
bench testing. Note that any system changing disinfectants is subject to the profiling and
benchmarking requirements as described in Section 1.3.1 and specified in 40 CFR § 141.172.
9.2 Pathogen Inactivation with Interactive Disinfectants
In 1988, several reports appeared on the combined efficiency of some disinfectants on pathogen
inactivation. Worley and Williams (1988) reported that a mixture of free chlorine and organic
N-halamine produced higher levels of inactivation of a variety of bacteria. The combination of free
chlorine and sodium bromide was also investigated and found to be more effective than using free
chlorine alone (Alleman et al., 1988). In a study at the University of Arizona, the synergistic
inactivation of E. coli and MS-2 coliphage was demonstrated by the combined application of
ehloramine and cupric chloride (Straub et al., 1994).
Recently there has been a great deal of interest in the potential of interactive disinfectants because
reports showed that some of these combinations are more effective for inactivating Cryptosporidium
(Finch et al., 1994). Research on interactive disinfectants for primary pathogen inactivation is under
way for several combinations of disinfectants:
• Chlorine followed by ehloramine;
• Chlorine dioxide followed by chlorine;
• Chlorine dioxide followed by chlorine dioxide;
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9. COMBINED DISINFECTANTS
• Chlorine dioxide followed by chloramine;
• Ozone followed by chlorine;
• Ozone followed by chlorine dioxide; and
• Ozone followed by chloramine.
9.2.1 Inactivation Mechanism
Bernbaum (1981 and 1985) developed a testing method for determining the kind of interaction that
can be expected when agents are combined to produce a given observation. Synergism can be tested
using the mathematical model developed by Bernbaum and modified for disinfection kinetics by
Kouame and Haas (1991). The principle is that, if the agents in a given combination do not interact
in producing the effect observed, then regardless of the effect relations, the following equation is
satisfied:
where:
Xj = Concentration of the individual agent in the combination
yt = Concentration of the agents that individually would produce the same magnitude
of effect as that of the combination
/ = Individual agent
n = Total number of agents
The sum calculated from this equation for a set of data is interpreted as follows:
• The sum is less than 1 in the case of synergistic interaction;
• The sum is greater than 1 in the case of antagonistic interaction; and
• The sum is equal to 1 in the case of additivity (zero interaction).
Using this approach, Kouame and Haas (1991) showed that a synergistic interaction exists in the
inactivation of E. coli when exposed simultaneously to free chlorine and monochloramine.
The authors described a possible mechanism in which both of the disinfectants work together to
inactivate bacteria. The researchers hypothesized that bacterial inactivation is caused by
monochloramine penetrating the cell and oxidizing thiol groups, which in turn causes structural
changes in the cell membrane. Once these changes have^been made, copper is allowed to pass into
the cell and binds either to sulfhydryl groups of respiratory enzymes or nucleic acids. More recently,
the researchers investigating E. coli inactivation hypothesized that a potential synergistic mechanism
consisting of sub-lethal injury caused by free chlorine resulted in enhanced sensitivity to
monochloramine (Kouame and Haas, 1991).
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Another hypothesis for the increased effectiveness of interactive disinfectants is that the first oxidant
(i.e., chlorine, chlorine dioxide, or ozone) conditions the outer membrane of Cryptosporidium
oocysts so that the secondary oxidant (i.e., chlorine, chlorine dioxide, and monochloramine) can
penetrate the oocyst more easily (Liyanage et al., 1996). For example, preliminary work on the
disinfection of Cryptosporidium parvum using free chlorine followed by monochloramine suggested
that there may be a synergism involving two chlorine species. Sequential treatment of these chlorine
species was found to provide greater inactivation than expected from the additive effects of the two
disinfectants used alone (Gyurek et al., 1996).
Recent studies have utilized a straight forward method to determine if synergism has occurred based
on measured inactivation (Finch, 1997; Gyurek et al., 1996; and Liyanage et al., 1996). According to
this approach, synergism is demonstrated if the sequential application of disinfectants provides more
inactivation than is expected from the additive effects of the individual, separate disinfectants. In
addition, the magnitude of the synergistic effects is equal to the difference in the level of inactivation
achieved from multiple disinfectants and the additive inactivations achieved from the single
disinfectants.
9.2.2 Environmental Effects
Similar to most chemical disinfectants, the preliminary results from an AWWARF ongoing study
suggest that pH and temperature affect the amount of synergistic inactivation achieved by sequential
applications of disinfectants (Finch, 1997). The following sections briefly describe the effects these
parameters have on pathogen inactivation.
9.2.2.1 pH
The level of inactivation due to the sequential application of chemical disinfectants is believed to be
pH dependent (Finch, 1997). Figure 9-1, Figure 9-2, and Figure 9-3 illustrate the impact of pH on
the log inactivation of Cryptosporidium parvum attributed to synergistic effects for three sequential
combinations of ozone-chlorine dioxide, chlorine dioxide-free chlorine' and chlorine dioxide-
chloramine, respectively. As shown in these figures, the amount of log inactivation due to
synergistic effects is lower at high pH (e.g., pH = 11). These results show that neutral pH is more
effective than low pH except for ozone-chlorine dioxide.
9.2.2.2 Combination of Low Temperature and pH
The combined effect of low temperature and high pH is believed to significantly reduce the amount
of Cryptosporidium inactivation attributed to synergism (Finch, 1997). One possible explanation for
this reduction is that the oocysts contract under these conditions and become harder to penetrate.
However, significant reduction in Cryptosporidium oocysts inactivation is true under reduced water
temperature and high pH whether interactive disinfection is practiced or not. Therefore, reduced
inactivation may not be necessarily due to synergism between combined disinfectants.
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9. COMBINED DISINFECTANTS
1.75-
1.50-
I
| 1.25-
| 1.00 -
1
| 0.75 •
0.25-
.. . .
urf
i^ it**
? tf
_
P#n
i^i
Ozone 0.9 mg/L for 4 minutes. Chlorine
dioxide 1 .3 mg/L for 120 minu es
6 8
PH
J
-------
9.
1
200 - • •
1.75
1.50
'E
3 1 25 -
0
o
'| 1.00-
c
_o
~ 0 75 -
o
n
0.25 -
0.00 -
181
Chlorine dioxide 1.3 mg/L for 120
minutes. Monochloramine: pH 6 and
11; 2.0 mg/L for 120 minutes; pH 8; 2.8
mg/L for 1 80 minutes.
6 8
pH
11
Figure 9-3. Inactivation of C. parvum Attributed to Synergistic Effects.
Application of Chlorine Dioxide Followed by Monochloramine
9.2.2.3 Pathogen Susceptibility
Cryptosporidium oocysts are more susceptible to inactivation by combinations of disinfectants than
by individual disinfectants. Giardia cysts were also found to have a similar response to:
• Ozone followed by free chlorine;
• Ozone followed by monochloramine;
• Chlorine dioxide followed by free chlorine;
• Chlorine dioxide followed by monochloramine; and
• Free chlorine followed by monochloramine.
However, no synergism was observed with bacterial spores, specifically Bacillus cereus spores
(Finch, 1997). These results suggest that encysted parasites might show more susceptibility to
synergistic effects than bacterial spores. Masking effects caused by turbidity for interactive
disinfectants are expected to be similar to those of the individual disinfectants.
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9. COMBINED DISINFECTANTS
9.2.3 Pathogen Iriactivation Efficiency Using Interactive
Disinfectants
Within the last few years there have been several studies to investigate interactive disinfectants.
These studies were conducted under various conditions of pH, bench or laboratory scale, and using
various organisms:
• Battigelli and Sobsey (1993) studied viral inactivation under lime softening conditions using
sequential addition of chlorine and monochloramine;
• Kouame and Haas (1991) determined E. coli inactivation during simultaneous addition of
free chlorine and monochloramine;
* Finch (1997) is studying various combinations of chlorine, chlorine dioxide, ozone, and
monochloramine on inactivation of Cryptosporidium parvum oocysts, Giardia muris cysts,
and Bacillus cereus spores under laboratory conditions; and
• Oppenheimer (1997) is developing CT requirements for Cryptosporidium parvum
inactivation in a variety of natural waters using ozone followed by chlorine.
The following is a summary of the findings of these studies to date.
9.2.3,1 Virus Inactivation Using Chlorine and Monochloramine Under High pH
One of the primary objectives of the Battigelli and Sobsey study (1993) was to evaluate the
disinfection efficiency under high pH conditions encountered in conventional lime softening
treatment with and without the addition of chlorine and monochloramine. The three microorganisms
selected for evaluation were hepatitis A virus, poliovirus 1, and MS-2 coliphage.
During the study, the inactivation kinetics of the three test viruses were determined when 2.0 mg/L
monochloramine were formed dynamically after the viruses had been exposed to lime solution and
free chlorine for 60 seconds. The authors believed that this approach simulates the conditions
typically encountered in a water softening facility where lammonia is applied post-chlorination.
Results indicated that a high degree of inactivation occurs during the first 60 seconds of chlorine
addition at approximately 2.4 mg/L free chlorine.
Table 9-6 shows the amount of inactivation attributed to the lime solution, free chlorine, and
monochloramine for the three viruses. The table also contains the amount of inactivation attributed
to the sequential application of lime solution, free chlorine and monochloramine previously described
(simultaneous chloramination). Results shown in Table 9-6 are based on a pH of 11.0 and a total
contact time of 360 minutes.
Except for poliovirus 1, the summation of the individual disinfectants was greater than the level of
inactivation achieved from the simultaneous chloramination. These results imply that the sequential
addition of free chlorine and monochloramine after lime addition to raise the pH to 11, form an
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9. COMSWED D(S(NFEGTANrS
antagonistic (negative) interaction for inactivation of hepatitis A virus and MS-2 coliphage. For
poliovirus 1, under similar conditions, an enhancement of 1.4 log inactivation was achieved
suggesting positive synergism for poliovirus 1 inactivation.
Table 9-6. Virus Inactivation By Individual Disinfectants and Simultaneous
Chloramination
Disinfectant(s)
Log Survival
Hepatitis Poliovirus MS-2
A Virus Coliphage
Condition
Lime solution only
Free chlorine
Monochloramine
Summation free+monochloramine
Simultaneous Chloramination
-3.0
-1.8
-3.7
-5.5
-4.5
-0.5
-1.2
-1.9
-3.1
-4.5
-4.0
-1.6
-3.8
-5.4
-4.5
360 min contact time
60 second contact time,
2.5 mg/L chlorine
2.0 mg/L monochloramine
Additive
2.4 mg/L chlorine 60
second contact time, 2.0
mg/L monochloramine,
359 minutes.
1 All data at pH 11 after lime addition
Source: Battigelli and Sobsey, 1993.
9.2.3.2 Inactivation of E. coll Under Simultaneous Free and Combined
Chlorination
The inactivation of E. coli bacteria by the simultaneous application of free chlorine and
monochloramine was investigated at the Illinois Institute of Technology (Kouame and Haas, 1991).
Figure 9-4 shows the level of E. coli inactivation by free chlorine and monochloramine, separately
and combined. The level of inactivation by monochloramine alone after a contact time of 300
seconds was found to be significantly less than that of free chlorine. Therefore, the sum of the
individual inactivation by free chlorine and monochloramine was assumed to be equal to that of free
chlorine alone. Note that in this case, the residual disinfectant rapidly disappeared due to the
breakpoint reactions that occur when monochloramine and free chlorine are combined.
The surviving fraction of bacteria following the simultaneous application of free chlorine and
monochloramine is substantially less than what would be expected by adding the individual levels of
inactivation. In other words, at similar doses and contact times, the amount of inactivation from the
combined disinfectants was greater than the sum of the inactivation due to free chlorine alone and
monochloramine alone.
In summary the Kouame and Haas (1991) study showed that high levels of bacteria inactivation can
be achieved when free chlorine and monochloramine exist simultaneously in a continuous flow
system and that the combined action of both chemicals on the bacteria is synergistic.
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9. COMBINED DISINFECTANTS
9.2.3.3 Cryptosporidium, Giardia, and Bacillus Inactivation in Laboratory Grade
Water
The preliminary results of an AWWARF study that investigated the application of multiple
disinfectants was presented at a American Water Works Association Technology Transfer
Conference in Portland, OR, in August 1997. The objectives of this study were to screen interactive
chemical disinfectants (ozone, chlorine, chlorine dioxide, and monochloramine) for inactivation of
Cryptosporidium parvum, Giardia muris, and Bacillus cereus and develop design criteria for
Cryptosporidium parvum inactivation using the best combinations.
Ozone Followed By Chlorine Dioxide
Table 9-7 shows the results obtained from ozone and chlorine dioxide application for the inactivation
of Cryptosporidium parvum.
\
Based on the data shown in Table 9-7, ozone followed by chlorine dioxide was the most effective
disinfectant combination for Cryptosporidium inactivation. A total contact time of 124 minutes was
required to achieve 3 to 4-log inactivation with ozone and chlorine dioxide residuals of 0.9 and 1.3
mg/L, respectively.
D Summation of expected levels of inactivation by free chlorine and monochloramine
• Levels of inactivation obtained from simultaneous application of free chlorine and monochloramine
Amount
attributed to
synergistic
effects. •
0.5
0.0
68 79
Contact Time (seconds)
Source: Kouame and Haas, 1991.
Figure 9-4. Inactivation of E. coll Using Free Chlorine and Monochloramine
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9. COMB/NED D(SINFECTANTS
Table 9-7. C. parvum Inactivation Using Ozone Followed by Chlorine Dioxide
Disinfectant
Ozone
Chlorine dioxide
Ozone followed by chlorine dioxide
Inactivation attributed to synergism
pHG.O
1.6
0.9
4.0
1.5
Level of Inactivation (log-units)
pH 8.0
0.8
1.4
3.6
1.4
pH11.0
0
2.4
2.9
0.5
Source: Finch, 1997.
Ozone: 0.9 mg/L for 4 minutes, chlorine dioxide 1.3 mg/L for 120 minutes.
Chlorine Dioxide Followed by Free Chlorine
Table 9-8 through Table 9-10 summarize of the results obtained for chlorine dioxide followed by free
chlorine for Cryptosporidium parvum, Giardia muris, and Bacillus cereus, respectively.
Table 9-8. C. parvum Inactivation Using Chlorine Dioxide Followed by Free
Chlorine
Disinfectant
Chlorine dioxide
Free chlorine
Chlorine dioxide followed by free chlorine
Inactivation attributed to synergism
pHG.O
1.0
0
2.2
1.2
Level of Inactivation (log-units)
pHS.O
1.4
0
3.0
1.6
pH11.0
1.6
0
2.3
0.7
Source: Finch, 1997.
Chlorine dioxide 1.3 mg/L for 120 minutes, free chlorine 2.0 mg/L for 120 minutes.
Table 9-9. G. muris Inactivation Using Chlorine Dioxide Followed by Free
Chlorine
Disinfectant Level of Inactivation (log-units)
pH6.0 pHS.O
Chlorine dioxide 0.8. 0.8
Free chlorine 0.8 0.6
Chlorine dioxide followed by free chlorine 2.2 2.0
Inactivation attributedI to synergism ——— — - __
Source: Finch, 1997. ———————___ _________
Chlorine dioxide: 1.0 mg/L for 10 minutes, free chlorine 2.0 mg/L for 30 minutes.
Table 9-10. B. cereus Inactivation Using Chlorine Dioxide Followed by Free
Chlorine
Disinfectant Level of Inactivation (log-units)
Chlorine dioxide 1~8 ~
Free chlorine 1.2
Chlorine dioxide followed by free chlorine 2.9
Inactivation attributed to synergism " 6
Source: Finch, 1997. ~ '
Chlorine dioxide: 2.3 mg/L for 20 minutes, free chlorine for 30 minutes.
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Chlorine dioxide followed by free chlorine was capable of achieving between 2 and 3 logs of
Cryptosporidium inactivation following a total contact time of 240 minutes and approximately 2 logs
of inactivation ofGiardia following only 40 minutes of contact time. No synergism was observed
with Bacillus cereus. However, approximately 3 logs of inactivation after a contact time of 50
minutes were achieved by the additive effects of chlorine dioxide and free chlorine.
Chlorine Dioxide Followed by Chioramine
Table 9-11 and Table 9-12 show the results of the inactivation of Cryptosporidium parvum and
Giardia muris when using chlorine dioxide followed by monochloramine.
Table 9-11. C. parvum Inactivation Using Chlorine Dioxide Followed by
Chioramine
Disinfectant
Chlorine dioxide
Monochloramine
Chlorine dioxide followed by monochloramine
Inactivation attributed to synergism
Level of Inactivation (log-units)
pHG.O pHS.Ci pH11.0
1.0
0
2.2
1.2
1.5
0
2.8
1.3
1.6
0
2.1
0.5
Source: Finch, 1997.
Chlorine dioxide: pH 6, 8, and 11:1.3 mg/L for 120 minutes. Monochloramine: pH 6 and 11: 2.0 mg/L for 120 minutes, pH 8:2.8
mg/L for 180 minutes.
Table 9-12. G. muris Inactivation Using Chlorine Dioxide Followed by Chioramine
Disinfectant Level of Inactivation (log-units)
pHS.O pH11.0
Chlorine dioxide '. 0.8 0.8
Monochloramine 0.5 0.7
Chlorine dioxide followed by monochloramine 1.7 1_.5
Inactivation attributed to synergism 0.4 0
Source: Finch, 1997. \.
pH 8.0: Chlorine dioxide 1.0 mg/L for 5 minutes, monochloramine 2.0 mg/L for 150 minutes.
pH 11.0: Chlorine dioxide 1.0 mg/L for 5 minutes, monochloramine 2.0 mg/L for 5 minutes.
At similar disinfect residuals and contact times, chlorine dioxide followed by monochloramine was
found to achieve the same levels of Cryptosporidium inactivation as chlorine dioxide followed by
free chlorine at pH values of 6 and 11. However, at pH 8, a higher monochloramine residual and
contact time were required to achieve inactivation levels comparable to chlorine dioxide and free
chlorine. No synergism was found for Giardia inactivation at pH 11.0 with only a minimal increase
in effectiveness at pH 8.0.
Ozone Followed by Free Chlorine
Table 9-13 and Table 9-14 show the levels of inactivation ofGiardia muris and Bacillus cereus
obtained by using ozone followed by free chlorine.
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9. COMBINED DISINFECTANTS
Table 9-13. G. muris inactivation Using Ozone Followed by Free Chlorine
Disinfectant
Ozone
Free chlorine
Ozone followed by free chlorine
Inactivation attributed to synergism
Level of Inactivation (log-units)
PH6.0 pH8.0 pH11.0
0.5
0.8
2.3
1.0
0.8
0.6
2.2
0.8
0.4
0
0.4
0
pH 6.0: ozone 0.1 mg/L for 60 seconds; free chlorine 2.0 mg/L for 30 minutes.
pH 8.0: ozone 0.1 mg/L for 17 seconds; free chlorine 2.0 mg/L for 60 minutes.
pH 11.0: ozone 0.1 mg/L for 5 seconds; free chlorine 2.0 mg/L for 60 minutes.
Table 9-14. B. cereus Inactivation Using Chlorine Dioxide Followed by Free
Chlorine
Disinfectant Level of Inactivation (log-units)
Chlorine dioxide ————-—- ^ 4 :
Free chlorine 2.0
Chlorine dioxide followed fay free chlorine 3.4
Inactivation attributed to synergism " o ~~~
Source: Finch; 1997. ' ' ' '
Ozone followed by free chlorine was capable of achieving approximately 2 logs of Cryptosporidium
inactivation at pH 6.0 and 8.0; however, only a 0.4 log inactivation was achieved at pH 11.0. The
difference in inactivations was primarily caused by the inability of free chlorine to inactivate
Cryptosporidium at pH 11.0. Similar to the disinfectant combination of chlorine dioxide and free
chlorine, no synergism was observed for Bacillus cereus inactivation; however, the additive effects
of ozone and free chlorine achieved greater than 3 logs of inactivation.
Ozone Followed by Monochloramine
Table 9-15 shows the results obtained for Giardia muris inactivation by ozone followed by
monochloramine.
Table 9-15. G. muris Inactivation Using Ozone Followed by Chloramine
Disinfectant Level of Inactivation (log-units)
- ; pHS.O BllJl?
Ozone .0.8 0^4
Monochloramine , 0.5 0.7
Ozone followed by monochloramine 2.1 1.8
Inactivation attributed to synergism 0.8 - rj.7
nch, 19977 ™— -..-. ____ - ;___ -
pH 8.0: ozone 0.1 mg/L for 17 seconds; monochloramine 2.0 mg/L for 150 minutes.
pH 1 1 .0: ozone 0.1 mg/L for 5 seconds; monochloramine 2.0 mg/L for 5 minutes.
Because of the different residuals and contact times, inactivation efficiencies of ozone followed by
chloramine and ozone followed by free chlorine could not be compared. However, for similar
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9. COMBINED DISINFECTANTS
monochloramine residuals, a very short monochloramine contact time of 5 minutes at pH 11 was
found to achieve a greater inactivation than a contact time of 150 minutes at pH 8.
Free Chlorine Followed by Monochloramine
Table 9-16 shows the results obtained for Giardia muris inactivation by free chlorine followed by
monochloramine.
Similar to the results obtained for ozone followed by monochloramine, a very short monochloramine
contact time of 5 minutes at pH 11 was found to achieve a greater inactivation than a contact time of
150 minutes at pH 8. However, free chlorine did not achieve any inactivation at pH 11.
Table 9-16. G. muris Inactivation by Free Chlorine Followed by Monochloramine
Disinfectant Level of Inactivation (log-units)
1 pHS.O pHH.0
Free chlorine °'6 °
Monochloramine O-5 Oi^
Free chlorine followed by monochloramine 2.4 OJ
Inactivation attributed to synergism 1-3 ° .,
Source: Finch, 1997.
pH 8.0: free chlorine 2.0 mg/L for 60 minutes; monochloramine 2.0 mg/L for 150 minutes.
pH 11.0: free chlorine 2.0 mg/L for 60 minutes, monochloramine 2.0 mg/L for 5 minutes.
9.2.3.4 Bench-Scale Tests Using Natural Waters
In another AWWARF study, Oppenheimer (1997) is developing CT requirements for
Cryptosporidium parvum inactivation in a variety of natural waters, developing design criteria for
full-scale contacting systems from bench scale CT values, and investigating the impact of selected
variables on CT requirements. To date, samples have been collected and analyzed from 13
geographically disperse locations. Although a significant amount of data were not available, results
from the California State Water Project and Ohio River appear to show that the sequential application
of ozone and chloramines resulted in an enhanced inactivation of C. parvum as shown in Table 9-17.
The sequential application of free chlorine and monochloramine appears to enhance C. parvum
inactivation by providing some synergistic effects. To obtain the log reduction, however, very high
ozone residuals were required which appear to be impractical. In addition, bromate formation was
also a problem.
Table 9-17. C. parvum Inactivation by Sequential Application of Ozone and
Chloramine
Ozone
Water Source
California State Water Project
Ohio River
Residual
mg/L
0.8
4
Contact
min
12
15
Chlorine
Residual
mg/L
1.5
1.5
Contact
min
~0
120
Chloramine
Residual
mg/L
2.5
0.5
Contact
mm
30
120
Log Inactivation
Enhancement
0.3 to > 1.4
0.9 to 1.4
Source: Oppenheimer, 1997.
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9. COMBINED DISINFECTANTS
9.2.4 Summary: Pathogen Inactivation with Interactive
Disinfectants
Various studies.have shown the synergistic effects of interactive disinfectants: either simultaneous
application or sequential application. The improved disinfection efficiency due to interactive
disinfection is variable, ranging from negative (antagonistic) effects (in two studies) to positive
enhancement of disinfection efficiency. Many of the studies show definite improvement in
inactivation for interactive disinfectants.
Several research projects on the effects of combined disinfectants are underway at the time this
manual is being prepared. These projects should provide insight on the mechanisms and applicability
of multiple disinfectants. Based on current information, EPA believes that under appropriate
situations a positive improvement in disinfection efficiency exists. This enhanced inactivation varies
from organism to organism, and with different disinfectant combinations. For the key organisms of
interest under normal pH conditions:
• Coliform bacteria inactivation appears to increase with combined disinfectants;
• Giardia cyst inactivation appears to increase with combined disinfectants;
• Hepatitis A virus and MS"2 coliphage inactivation using combined disinfectants appears to be
less efficient than the individual disinfectants;
• Poliovirus 1 inactivation appears to increase with combined disinfectants;
• Cryptosporidium oocyst inactivation appears to increase with combined disinfectants; and
• Inactivation of spores appears neutral.
Interactive disinfection is still however considered an emerging technology. As such, CT credits for
interactive disinfectants have not yet been established.
9.3 Analytical Methods
In general, most of the analytical methods for residual disinfectants are impacted negatively by the
presence of other disinfectants. Fortunately, for most of the disinfectants and oxidants listed below, at
least one method exists that can be used successfully in the presence of other oxidizing agents. For
analytical method details, see the individual disinfectant chapters.
9.3.1 Ozone
Residual ozone analysis cannot be performed in the presence of other oxidizing agents including
chlorine, chloramine, and potassium permanganate. Typically, the ozone analytical methods exhibit
interferences from chlorine, bromine, iodine, and manganese ions. The ACVK is the least susceptible
to interference and can be used when manganese concentrations are less than 1 mg/L and free or
combined chlorine concentrations are less than 10 mg/L (Gordon et al., 1992).
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9. COMBINED DISINFECTANTS
9.3.2 Chlorine Dioxide
Some of the analytical methods for chlorine dioxide, chlorate, and chlorite cannot be performed in
the presence of oxidizing agents. Amperometric and iodometric methods cannot be used in the
presence of metal ions such as manganese. Analytical methods that can be used in the presence of
other disinfectants and oxidants include UV spectrophotometric methods and ion chromatography
(Gordon et al., 1992).
9.3.3 Potassium Permanganate
The atomic adsorption method for permanganate analysis can be performed in the presence of any of
the other disinfectants (Standard Methods, 1995).
9.3.4 Chloramine
None of the colorimetric analytical methods for chloramine can be performed in the presence of
oxidizing agents such as ozone or hydrogen peroxide. Analytical methods that can be used in the
presence of other disinfectants and oxidants include the UV spectrophotometric method and the
amperometric titration methods (Gordon et al., 1992).
9.3.5 Hydrogen Peroxide
The analytical procedures for hydrogen peroxide in drinking water are all impacted by other
oxidizing species such as ozone and chlorine (Gordon et al., 1992).
9.3.6 UV Radiation
There are no known interferences from other disinfectants with the measurement of UV radiation
(DeMers and Renner, 1992). ',
l
9.3.7 Summary of Analytical Methods
Ozone analysis in the presence of chlorine is limited. However, these disinfectants are not commonly
present simultaneously, especially with the rapid decomposition of ozone.
Hydrogen peroxide analysis is difficult in the presence of ozone and other oxidizing agents.
However, when using peroxone, the ozone residual is the analyte used to meet disinfection
requirements.
All of the other disinfectants and oxidizing agents can be selectively monitored in the presence of
other disinfectants.
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9. COMBINED DISINFECTANTS
9.4 Summary
Table 9-18 summarizes, in general, the factors and uses of combined disinfectants. Specific
considerations depend on the actual combination of disinfectants used.
Table 9-18. Summary of Combined Disinfectants
Consideration
Generation
Primary uses
Inactivation efficiency
Byproduct formation
Limitations
Point of application
Special considerations
Description
Generation depends on the type of chemicals used. Ozone,
chlorine dioxide, and chloramines require on-site generation.
Two separate disinfectants can be used to provide primary
and secondary disinfection. By separating the primary and
secondary disinfection functions, the processes can be
optimized for maximum inactivation and minimum DBP
formation.
Interactive disinfection (using synergism between two
disinfectants to enhance inactivation) can serve as a primary
disinfectant.
The use of interactive disinfection as primary disinfectant for
inactivation of Giardia, Cryptosporidium, and viruses are
feasible. Interactive disinfection is typically more effective
than the individual disinfectants.
DBP formation is in general reduced by using combined
disinfectants. Specifically, continued use of chlorine in
combination with other disinfectants can reduce DBP
formation.
Data on the inactivation efficiency of combined disinfectants
are still being generated with much information coming from
controlled laboratory studies. Additional information is still
needed, specifically on full-scale implementation. Dual
(primary/secondary) disinfection for DBP control is well
established as a preferred treatment option.
Applied for primary and secondary disinfection. Ozonation
should occur after settling and prior to biofiltration.
The efficiency and application of combined disinfectants
follow to a large extent the limitations and features of the
individual disinfectant. The combined disinfectant is often a
more effective disinfectant.
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9.5 References
1. Alleman, J. E., et al. 1988. "Comparative evaluation of alternative halogen-based disinfection
strategies." Conference proceedings, Industrial Waste Conference, forty-second edition.
2. Battigelli, D.A. and M.D. Sobsey. 1993. "The Inactivation of Hepatitis A Virus, Poliovirus and
Coliphage MS2 By Lime Softening and Chlorine/Monochloramine Disinfection." Conference
proceedings, AWWA Water Quality Technology Conference.
3. Bernbaum, C.M. 1981. "Criteria for Analyzing Interactions Between Biologically Active
Agents." Adv. Cancer Res. 35:269.
4. Bernbaum, C.M. 1985. "The expected effect of a combination of agents: the general solution."
J.Theor.Biol. 114:413.
5. DeMers, L.D. and R.C. Renner. 1992. Alternative Disinfection Technologies for Small Drinking
Water Systems. AWWA and AWWARF, Denver, CO.
6. Finch, G.R. 1997. "Control of Cryptosporidium Through Chemical Disinfection: Current State-
of-the-Art." AWWARF Technology Transfer Conference, Portland, Oregon.
7. Finch, G.R., E.K. Black, and L.L. Gyurek. 1994. "Ozone and Chlorine Inactivation of
Cryptosporidium." Conference proceedings, Water Quality Technology Conference; San
Francisco, CA.
8. Gordon, G., W.J. Cooper, R.G. Rice, and G.E. Pacey. 1992. Disinfectant Residual Measurement
Methods. Second Edition, AWWARF and AWWA, Denver, CO.
9. Gyurek, L., L. Liyanage, M. Belosevic, and G. Finch, 1996. "Disinfection of Cryptosporidium
Parvum Using Single and Sequential Application of Ozone and Chlorine Species." Conference
proceedings, AWWA Water Quality Technology Conference, Boston, MA.
j
10. Hubbs, S., D. Amundsen, and P. Olthius. 1981. "Use of Chlorine Dioxide, Chloramines, and
Short-Term Free Chlorination as Alternative Disinfectants." /. AWWA. 73(2):97-101.
11. Jacangelo, J.G., N.L. Patania, K.M. Reagan, E.M. Aieta, S.W. Krasner, and M.J. McGuire. 1989.
"Impact of Ozonation on the Formation and Control of Disinfection Byproducts in Drinking
Water," J. AWWA. 81 (8):74. !
12. Kouame, Y. and C.N. Haas. 1991. "Inactivation of E coli. by Combined Action of Free Chlorine
and Monochloramine." Water Res. 25(9): 1027. ;
13. LeLacheur, R.M., P.C. Singer, and M.J. Charles. 1991. "Disinfection Byproducts in New Jersey
Drinking Waters." Conference proceedings, AWWA Annual Conference, Philadelphia, PA.
EPA Guidance Manual 9-26 \ April 1999
Alternative Disinfectants and Oxidants
-------
9. COMBINED DISINFECTANTS
14. Liyanage, L., L. Gyurek, M. Belosevic, and G. Finch. 1996. "Effect of Chlorine Dioxide
Preconditioning on Inactivation of Cryptosporidium by Free Chlorine and Monochloramine."
Conference proceedings, AWWA Water Quality Technology Conference, Boston, MA.
15. Logsdon, G.S., S. Foellmi, and B. Long. 1992. Filtration Pilot Plant Studies for Greater
Vancouver's Water Supply. Conference proceedings, AWWA Annual Conference, Toronto,
Ontario.
16. Lykins, B.W., J.A. Goodrich, W.E. Koffskey, and M.H. Griese. 1991. "Controlling Disinfection
Byproducts with Alternative Disinfectants." Conference proceedings, AWWA Annual
Conference, Philadelphia, PA.
17. Malcolm Pirnie, Inc. 1990. "Task 1.3 - Water Quality Report." Prepared for the Public Utilities
Department, City of San Diego.
18. Malcolm Pirnie, Inc. 1989. "Water Quality Master Plan." City of Phoenix Water and Wastewater
Department.
19. McKnight, A. and D. Reckhow. 1992. "Reactions of Ozonation Byproducts with Chlorine and
Chloramines." Conference proceedings, AWWA Annual Conference, Vancouver, British
Columbia.
20. MWDSC AND JMM (Metropolitan Water District of Southern California and James M
Montgomery Consulting Engineers). 1989. Disinfection Byproducts in United States Drinking
Waters. Volume I. EPA and Association of Metropolitan Water Agencies. Cincinnati, OH and
Washington, D.C.
21. Oppenheimer, J.A. 1997. "Cryptosporidium Inactivation in Natural Waters." AWWARF
Technology Transfer Conference, Portland, OR.
22. Reckhow, D., B. Legube, and P. Singer. 1986. "The Ozonation of Organic Halide Precursors:
Effect of Bicarbonate." Water Res. 20(8):987-998.
23. Richardson, Susan D., Alfred D. Thurston, Timothy W.Collette, Kathleen Schenck Patterson,
Benjamin w. Lykins, George Majetich, and Yung Zhang. 1994. Multispecial Identification of
Chlorine Dioxide Disinfection Byproducts in Drinking Water. Environ. Sci. Technol. 28:4:592.
24. Singer, P.C. 1988. Alternative Oxidant and Disinfectant Treatment Strategies for Controlling
Trihalomethane Formation. EPA Risk Reduction Engineering Laboratory, Cincinnati OH Rept.
No. EPA/600/2-88/044.
25. Singer, P.C. 1992. Formation and Characterization of Disinfection Byproducts. Presented at the
First International Conference on the Safety of Water Disinfection: Balancing Chemical and
Microbial Risks.
April 1999 9-27 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
9. COMBINED DISINFECTANTS
26. Standard Methods. 1995. Standard Methods for the Examination of Water and Wastewater,
nineteenth edition. Franson, M.H., Eaton, A.D., Clesceri, L.S., and Greenberg, A.E., (editors),
American Public Health Association, AWWA, and Water Environment Federation, Washington
D.C..
27. Straub, T. M., et al. 1994. "Synergistic Inactivation of Escherichia coli and MS-2 Coliphage by
Chloramine and Cupric Chloride." Conference proceedings, AWWA Water Quality Technology
Conference, San Francisco, CA.
28. USEPA (U.S. Environmental Protection Agency). 1992. Technologies and Costs for Control of
Disinfection Byproducts. Prepared by Malcolm Pirnie, Inc. for the Office of Ground Water and
Drinking Water, Report No. PB93-162998. '
29. Worley, S.D. and D.E. Williams. 1988. "Disinfecting Water with a Mixture of Free Chlorine and
Organic B-halamine." J. AWWA. 80(1):6
EPA Guidance Manual 9-28 ' April 1999
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT
USAGE IN THE UNITED STATES
Two sources of information were consulted regarding disinfectant usage in the United States:
• The Community Water Systems Survey (USEPA, 1997); and
• The Information Collection Rule (ICR) database on water utilities (presently under
development).
A.1 Community Water Systems Survey
Most water treatment plants disinfect water prior to distribution. The 1995 Community Water
Systems Survey (USEPA 1997a) reports that 81 percent of all community water systems provide
some form of treatment on all or a portion of their water sources (Table A-l). The survey found
that 99 percent of the surface water systems provide some treatment of their water. Of those
systems reporting no treatment, 80 percent rely on ground water as their only water source.
Table A-1. Disinfection Practices of Water Systems with Treatment
Treatment
<100
101-500 501-1,000
Service Population
1,001- 3,301-
3,300 10,000
10,001-
50,000
50,001-
100,000
Over
100,001
Pre-Disinfection, Oxidation/Softening
Chlorine 59.0% 73.9%
Chlorine Dioxide 0.0% 0.0%
Chloramines 4.6% 0.0%
Ozone 0.0% 0.0%
KMn04 0.0% 4.9%
Predisinfection/oxidation 0.0% 0.0%
Lime/Soda ash softening 6.8% 9.8%
Recarbonation 0.0% 0.0%
Post-Disinfection
Chlorine 49.7% 51.6%
Chlorine Dioxide 0.0% 0.0%
Chloramines 0.0% 0.0%
Postdisinfection combine 0.0% 0.0%
Fluoridation 0.0% 4.9%
Pre-Disinfection, Oxidation/Softening
Surface Water Systems
67.3%
0.0%
1.1%
0.0%
9.6%
2.0%
20.9%
0.0%
80.6%
0.0%
0.0%
0.0%
13.9%
66.3%
5.0%
2.1%
0.0%
9.9%
2.9%
16.2%
0.0%
62.8%
0.0%
2.9%
2.1%
32.4%
68.8%
4.7%
0.0%
0.3%
15.2%
0.6%
14.3%
2.1%
77.9%
0.3%
2.1%
4.0%
42.6%
58.6%
13.2%
2.2%
0.0%
28.3%
9.2%
11.7%
4.7%
71.1%
4.9%
15.6%
3.9%
48.8%
47.5%
14.2%
15.5%
5.4%
25.9%
5.1%
3.5%
0.6 %
73.8%
5.9%
29.4%
1.9%
49.9%
Ground Water Systems
Source: 1995 Community Water Systems survey (USEPA, 1997a)
57.1%
7.8%
10.8%
5.8%
28.5%
4.3%
5.9%
6.3%
63.6%
11.2%
24.2%
1.6%
63.6%
Total
63.8%
6.3%
3.1%
0.9%
16.0%
3.5%
12.5%
1.9%
67.5%
1.6%
8.1%
3.0%
35.5%
Chlorine
Chlorine Dioxide
Chloramines
Ozone
KMn04
Predisinfection/oxidation
Lime/Soda ash softening
Recarbonation
Post-Disinfection
Chlorine
Chlorine Dioxide
Chloramines
Postdisinfection combine
Fluoridation
64.2%
1.3%
0.0%
0.0%
0.0%
0.3%
2.9%
0.0%
23.0%
0.0%
' 0.0%
0.0%
2.4%
69.9%
0.0%
0.0%
0.0%
0.9%
0.5%
2.9%
0.5%
23.4%
1.0%
0.0%
0.0%
6.3%
56.7%
0.0%
0.0%
0.0%
2.2%
0.0%
2.2%
0.0%
32.5%
0.0%
0.0%
0.0%
13.2%
73.2%
0.0%
0.0%
0.0%
0.6%
0.7%
3.6%
0.6%
28.3%
0.0%
0.0%
0.0%
12.4%
60.6%
0.0%
0.0%
0.0%
5.8%
1.0%
3.5%
1.4%
42.5%
0.0%
0.1%
0.1%
45.3%
57.4%
0.0%
0.6%
0.0%
3.2%
2.6%
3.8%
1.5%
41.9%
0.6%
1.1%
0.1%
31.2%
36.2%
3.1%
1.4%
0.0%
7.0%
0.0%
5.0%
2.8%
54.5%
0.0%
3.9%
,0.0%
34.3%
38.1%
0.0%
0.7%
0.6%
0.0%
0.0%
9.1%
1.1%
65.8%
0.0%
4.3%
0.0%
52.5%
63.9%
0.3%
0.1%
0.0%
1.8%
0.7%
3.2%
0.6%
31.0%
0.4%
0.3%
0.0%
16.0%
April 1999
A-1
EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT USAGE IN THE U. S.
A.2 Information Collection Rule Database
The ICR database contains information from 527 community water systems. Some of these
systems are owned by the same utility, representing different treatment plants within the same
organization. The following tables summarize the ICR data regarding disinfectant database
entries. Note that the information about disinfectant includes usage of the disinfectant at the
given utility. In some cases, the utility may use more than one disinfectant. Therefore, the 527
systems reported use of 740 different disinfectants. Three systems reported using 3 different
chemical disinfectants - for different purposes or during different times of the year. The tables
show a breakdown of the disinfectant usage at these facilities. Capacities are shown in terms of
flow. Population served data are not recorded.
The tables are as follows:
• Table A-2 shows the breakdown of systems based on water source (i.e., surface or ground
water).
• Table A-3 to Table A-5 shows the disinfectant usage at water plants by flow categories for
all water sources, surface water sources, and ground water sources, respectively. The tables
show the percentage of facilities that are using the particular disinfectant. Because some
facilities use more than one disinfectant, the total usage exceeds 100 percent. For example,
for facilities in the range 51-100 mgd, the total disinfectant usage is 155 percent. This means
that 155 types of disinfection systems are used at 100 plants.
• Table A-6 through Table A-8 shows the disinfectant usage in number of applications in water
plants by flow categories for all water sources, surface water sources, and ground water
sources, respectively. These tables show the actual numbers used in calculating the
percentages in Table A-3 to Table A-5.
• Table A-9 through Table A-l 1 show the number of systems using two or more disinfectants
for all water sources, surface water sources, and ground water sources, respectively. The
database does not separate usage as primary or secondary disinfectant.
Table A-2. Breakdown of systems in Survey based on Water Source
, II-IITII-IIIIIH-T-- j_ijij...jiLijiiimrrnmr-|-r-|.imr.miiriiiitiiii r:|-|iiinnirTi mmmimim I.iiTTiiiuiiiui...m...i..T.r,rmriiimiiliitiiititttnmiiinlililililtnnmmimnilmrnnlrmrtiimrrlr
Systems GW SW GW/SW Unknown Total
Number 135 390 2 2 529
Percentage 26% 74% 0% 0% 100%
GW = Ground Water; SW = Surface Water; GW/SW = Choice of Ground or Surface Water Source.
EPA Guidance Manual A-2 April 1999
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT USAGE IN THE U. S.
Table A-3. Disinfectant usage as a function of flow for all Water Sources.
Numbers show the percentage of systems using a particular disinfectant
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Percentage*
CI2
69%
90%
93%
95%
98%
77%
92%
NaOCI
19%
10%
6%
5%
5%
13%
7%
NH2CI2
19%
10%
30%
41%
43%
8%
31%
03
0%
0%
5%
5%
5%
0%
4%
CIO2
, 0%
0%
9%
9%
3%
0%
6%
Total use
106%
110%
143%
155%
154%
98%
140%
Percentage calculated as a fraction of 527 - the total number of systems. 740 different disinfectants
are used by the 527 systems.
Table A-4. Disinfectant usage as a function of flow for Surface Water Sources.
Numbers show the percentage of systems using a particular disinfectant
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Percentage*
CI2
89%
85%
93%
96%
96%
88%
94%
NaOCI
11%
8%
5%
4%
6%
13%
5%
NH2CI
0%
8%
32%
42%
46%
50%
37%
03
0%
0%
5%
6%
6%
0%
5%
CIO2
0%
0% x
12%
11%
4%
0%
9%
Total use
100%
100%
148%
158%
158%
150%
150%
Percentage calculated as a fraction of 527 - the total number of systems. 576 different disinfectants
are used by the 383 systems.
Table A-5. Disinfectant usage as a function of flow for Ground Water Sources.
Numbers show the percentage of systems using a particular disinfectant
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Percentage*
CI2
43%
94%
92%
88%
108%
79%
87%
NaOCI
29%
13%
6%
13%
0%
14%
10%
NH2CI
43%
13%
24%
38%
25%
0%
18%
03
0%
0%
2%
0%
0%
0%
1%
CIO2
0%
0%
2%
0%
0%
0%
1%
Total use
114%
119%
125%
138%
133%
93%
117%
Percentage calculated as a fraction of 527 - the total number of systems. 168 different disinfectants
are used by the 144 systems.
April 1999 A-3 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT USAGE IN THE U.S.
Table A-6. Disinfectant usage as a function of flow for all Water Sources
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Total
Percentage*
CI2
11
26
200
113
94
40
484
92%
NaOCI
3
3
12
6
5
7
36
7%
NH2CI
3
3
65
49
41
4
165
31%
03
0
0
10
6
5
0
21
4%
CIO2
0
0
20
11
3
0
34
6%
Total use
17
32
307
185
148
51
740
140%
Total
plants
16
29
215
119
96
52
527
* Percentage calculated as a fraction of 527 - the total number of systems. 740 different disinfectants
are used by the 527 systems.
Table A-7. Disinfectant usage as a function of flow for Surface Water Sources
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Total
Percentage*
CI2
8
11
154
99
82
7
361
94%
NaOCI
1
1
9
4
5
1
21
5%
NH2CI
0
1
53
43
39
4
140
37%
03
0
0
9
6
5
o :
20
5%
CIO2
0
0
20
11
3
0
34
9%
Total use
9
13
245
163
134
12
576
150
Total
plants
9
13
165
103
85
8
383
* Percentage calculated as a fraction of 383 - the total number of plants. 576 different disinfectants
are used by the 383 systems.
I
Table A-8. Disinfectant usage as a function of flow for Ground Water Sources
Flow, mgd
0-5
6-10
11-50
51-100
>100
Unknown
Total
Percentage*
CI2
3
15
47
14
13
33
125
87%
NaOCI
2
2
3
2
0
6
15
10%
NH2CI
3
2
12
6
3
0
26
18%
03
!
0
0
1 ;
0
0
0
1
1%
CIO2
0
0
1
0
0
0
1
1%
Total use
8
19
64
22
16
39
168
117%
Total
plants
7
16
51
16
12
42
144
' Percentage calculated as a fraction of 144 - the total number of systems. 168 different disinfectants
are used by the 144 systems.
EPA Guidance Manual A-4 April 1999
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT USAGE IN THE U.S.
Table A-9. Number Water Systems (Ground and Surface Water Sources)
using Two Different Disinfectants
CI2
NaOCI
NH2CI
03
CI02
CI2 NaOCI
8
—
—
—
—
NH2CI
150
8
—
—
—
03
17
3
12
—
—
CIO2
32
0
18
1
—
Table A-10. Number Surface Water Systems using Two Different Disinfectants
CI2
NaOCI
NH2CI
03
CI02
CI2 NaOCI NH2CI
7 129
7
...
—
—
03
16
7
11
—
—
CI02
31
3
18
1
—
Table A-11. Number Ground Water Systems using Two Different Disinfectants
CI2~ NaOCI NH2CI~ O3 CIO2
_ ^ 1 21 1 0
NaOCI — — 1 0 0
NH2CI — — — 1 0
03 — — — — 0
CIO2
April 1999 A-5 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
APPENDIX A - SUMMARY OF DISINFECTANT USAGE IN THE U. S.
THIS PAGE INTENTIONALLY LEFT BLANK
EPA Guidance Manual
Alternative Disinfectants and Oxidants
A-6
April 1999
-------
Appendix B - Selected Costs of
Alternative Disinfection Systems
B.1 Technologies and Costs for Control of
Disinfection By-Products
Costs were developed for modifying a "base" or "typical" treatment plant to add disinfection and
other technologies. The base plant is described as a conventional treatment plant using
chlorine/chlorine disinfection consisting of rapid mixing, flocculation, sedimentation, chlorination,
filtration, contact basin, chemical feed systems and finished water storage. This appendix contains
figures and tables from Technologies and Costs for Control of Disinfection By-Products (USEPA,
1998), retaining the report's original figure and table numbers. Incremental costs are shown,
determined by calculating the cost for the modified treatment plant and subtracting the base
treatment plant cost.
The base treatment plant shown in Figure 7-1, is a basic alum coagulation and filtration plant, with
chlorine disinfection. This plant was modified to meet disinfection requirements. The bases for the
cost estimates are shown in Tables 7-3, 7-4, 7-5, and 7-6. The 12 flow categories for which the costs
were determined are shown in Table 7-2.
Schematics and costs to add the following schemes are shown in the attached figures and tables.
• Base treatment plant - Figure 7-1 and Table 7-7.
• Move point of chlorination. This modification assumes no cost for moving the chlorine addition
point, but costs for an added contact basin are shown in Table 7-8.
• Change to Chlorine/Chloramine - Figure 7-2 and Table 7-9.
• Change to Ozone/Chloramine - Figure 7-3 and Table 7-12.
• Change to Chlorine Dioxide - Table 7-13.
See USEPA, 1998 for more details and information upon the costs of other technologies.
April 1999 B-1 EPA Guidance Manual
Alternative Disinfectants and Oxidants
-------
APPENDIX B - COST OF ALTERNATIVE DISINFECTION SYSTEMS
FIGURE 7-1. ALUM COAGULATION / FILTRATION BASE PLANT
•Alum
Rapid
Mix
CNortne
• Caustic
Rocculation
& Clarification
Filtration Contact
Basin
Storage
TABLE 7-2. EPA FLOW CATEGORIES
EPA FLOW
CATEGORIES
MEDIAN
POPULATION
SERVED
AVERAGE FLOW
(mgd)
DESIGN
CAPACITY
(mgd)
SMALL SYSTEMS - DESIGN FLOW < 1 MGD
1
2
3
4
57
225
750
1,910
0.0056
0.024
0.086
0.23
0.024
0.087
0.27
0.65
LARGE SYSTEMS - DESIGN FLOW > 1 MGD
5
6
7
8
9
10
11
12
12a
5,500
15,000
35,000
60,000
88,000
175,000
730,000
1,550,000
N/A
0.70
2.1
5.0
8.8
13
27
120
270
350
1.8
4.8
11
18
26
51
210
430
520
EPA Guidance Manual
Alternative Disinfectants and Oxidants
B-2
April 1999
-------
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-------
APPENDIX B - COST OF ALTERNATIVE DISINFECTION SYSTEMS
TABLE 7-5. COST ALLOWANCE FACTORS
Item
Site work and Interface Piping
Subsurface Considerations
Standby Power
General Contractors Overhead
and Profit
Engineering
Legal, Fiscal and
Administration fees
Small Systems
Water Model (%)
10
10
5
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5 to 6C
CI)
Large Systems WATERCO$T
Model(%)
15
10
5
12
<2)
15
9toll<2)
Notes:
"'Percentages added to estimated construction cost plus estimated cost for other allowances
factors.
<2>
>Percentages added to estimated construction cost only.
TABLE 7-6. INDICES USED IN THE ESCALATION OF COSTS
DESCRIPTION
Building Cost Index
Chemical & Allied Products
Skilled Labor
Materials
Utility Natural Gas
INDEX
REFERENCE
ENR1
BLS2
ENR1
ENR1
BLS 055 2
NUMERICAL
VALUE
3391.86
147.2
5231.35
2268.57
111.3
ESCALATION
VALUE
1.23
1.19
1.178
1.328
1.679
Engineering News Record (July, 1997)
; Bureau of Labor Statistics (March, 1997)
EPA Guidance Manual
Alternative Disinfectants and Oxidants
B-8
April 1999
-------
APPENDIX B - COST OF ALTERNATIVE DISINFECTION SYSTEMS
TABLE 7-7. ESTIMATED BASE PLANT COSTS
SMALL SYSTEMS
Design
Flow
(mgd)
0.024
0.087
0.27
0.65
Capital
Cost1
($M)
0.63
0.86
1.4
2.0
O&M
Cost2
(0/1000 gal)
600
188
90
56
Total
Cost @ 3%
(0/1000 gal)
2672
848
390
216
Total
Cost @ 7%
(0/1000 gal)
3509
1115
496
277
Total
Cost @ 10%
(0/1000 gal)
4233
1343
605
330
LARGE SYSTEMS
Design
Flow
(mgd)
1.8
4.8
11
18
26
51
210
430
520
Capital
Cost1
($M)
4.3
7.3
12
17
22
36 .
120
230
380
O&M
Cost2
(0/1000 gal)
74
47
39
36
35
33
32
31
26
Total
Cost @ 3%
(0/1000 gal)
187
111
83
72
66
58
50
47
46
Total
Cost @ 7%
(0/1000 gal)
233
137
102
86
78
68
58
53
54
Total
Cost @ 10%
(0/1000 gal)
272
159
118
98
89
76
65
59
61
1 1991 Cost escalated based upon a factor of 1.23 derived from the ENR BCI
2 1991 Cost escalated based upon a factor of 1.19 derived from the BLS Chemical and
Allied Products Index
April 1999
B-9
EPA Guidance Manual
Alternative Disinfectants and
-------
APPENDIXB - COST OF ALTERNATIVE DISINFECTION SYSTEMS
TABLE 7-8. ESTIMATED UPGRADE COSTS FOR ADDITIONAL CONTACT
BASIN SIZE (x $1000)*
DESIGN
FLOW
0.024
0.087
0.27
0.65
1.8
4.8
11
18
26
51
210
430
520
Chlorine Contact Basin Time
30min
14
25
52
77
197
274
432
611
815
1,454
5514
11,132
13,374
60 min
21
34
80
112
244
396
713
1,070
1,478,
2,755
10,876
22,112
26,639
120 min
26
66
103
218
335
642
1,274
1,990
2,807
5,360
21,600
44,071
53,224
180 min
28
76
140
251
427
887
1,836
2,909
4,135
7,965
32,324
66,031
79,785
240 min
38
82
180
284
519
1,132
2,399
3,828
5,462
10,569
43,050
87,991
106,352
300 min
46
84
220
317
611
1,376
2,961
4,748
6,791
13,175
53,774
109,951
132,193
360 mim
55
100
234
351
702
1622
3521
5667
8118
15,778
64,499
131,910
159,456
1 1991 Cost escalated based upon a factor of 1.23 derived from the ENR BCI
EPA Guidance Manual
Alternative Disinfectants and Oxidants
B-10
April 1999
-------
APPENDIX B - COST OF ALTERNATIVE DISINFECTION SYSTEMS
FIGURE 7-2. ALUM COAGULATION / FILTRATION SYSTEM UPGRADED WITH
CHLORINE / CHLORAMINE DISINFECTION
p Alum
Rapid Rocculafion
Mix & Clarification
Filtration
Contact
•Basin
Caustic
Ammonia
Storage
CHLORAMINATION PROCEDURE
Ammonia Dose Based on 4:1 Chlorine Residual to Ammonia Ratio
TABLE 7-9. ESTIMATED UPGRADE COSTS FOR CHLORAMINES AS SECONDARY
DISINFECTANT
SMALL SYSTEMS
Design
Flow
(mgd)
0.024
o:o8?
0.27
0.65
Capital
Cost1
($M)
0.011
0.012
0.015
0.016
O&M
Cost2
(0/1000 gal)
21
5.5
1.9
0.98
Total
Cost @ 3%
(0/1000 gal)
57
15
5.1
2.3
Total
Cost @ 7%
(0/1000 gal)
71
19
6.3
2.8
Total
Cost @ 10%
(0/1000 gal)
83
1 22
7.4
3.2
LARGE SYSTEMS
1991 Cost escalated based upon a factor of 1.23 derived from the ENR BCI
2 1991 Cost escalated based upon a factor of 1.19 derived from the BLS Chemical and
Allied Products Index
Design
Flow
(mgd)
1.8
4.8
11
18
26
51
210
430
520
Capital
Cost1
($M)
0.04
0.07
0.11
0.16
0.21
0.28
0.47
0.85
0.91
O&M
Cost2
(0/1000 gal)
1.4
0.70
0.49
0.40
0.37
0.33
0.29
0.26
0.20
Total
Cost @ 3%
(0/1000 gal)
2.5
1.3
0.9
0.73
0.67
0.52
0.36
0.32
0.25
Total
Cost @ 7%
(0/1000 gal)
3.0
1.5
1.1
0.87
0.79
0.60
0.39
0.34
0.27
Total
Cost @ 10%
(0/1000 gal)
3.4
1.8
1.2
0.99
0.89
0.67
0.41
0.36
0.28
April 1999
B-11
EPA Guidance Manual
Alternative Disinfectants and
-------
APPENDIXB- COST OF ALTERNATIVE DISINFECTION SYSTEMS
FIGURE 7-3. ALUM COAGULATION / FILTRATION SYSTEMS UPGRADED WITH
OZONE / CHLORAMINE DISINFECTION
Rapid Flocculation
Mix & Clarification
Ozone
Caustic
Chlorine
Ammonia
Filtration
Storage
CHLORAMINAT1ON PROCEDURE
Free Chlorine Contact for 1 Minute at
Peak Hourly Flow Prior to Ammonia Addition
Ammonia Dose Based on 4:1 Chlorine Residual to Ammonia Ratio
EPA Guidance Manual
Alternative Disinfectants and Oxidants
B-12
April 1999
-------
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APPENDIX B - COST OF ALTERNATIVE DISINFECTION SYSTEMS
B.2 References
1. USEPA. 1998a. Technologies and Costs for Control of Disinfection By-Products. Washington
DC. |
2. USEPA. 1998b. Regulatory Impact Analysis for the Stage 1 Disinfectant/Disinfection Byproduct
Rule. Prepared by Science Applications International Corporation for the USEPA, Office of
Ground Water and Drinking Water, Washington, DC.
AU.S. GOVERNMENT PRINTING OFFICE: 1999-45 2-143/10340
April 1999
B-28
EPA Guidance Manual
Alternative Disinfectants and Oxidants
------- |