MIAMI UNIVERSITY The Graduate School Certificate for

MIAMI UNIVERSITY

The Graduate School

Certificate for Approving the Dissertation

We hereby approve the Dissertation

of

Aleksey N. Pisarenko

Candidate for the Degree:

Doctor of Philosophy

_____________________________________ Director

Dr. Gilbert E. Pacey

_____________________________________

Director (Committee Chairperson) Dr. Gilbert Gordon

_____________________________________

Reader Dr. Richard T. Taylor

_____________________________________

Reader Dr. Michael W. Crowder

_____________________________________

Graduate School Representative Dr. Luis A. Actis

Abstract

ANALYTICAL MEASUREMENTS AND PREDICTIONS OF PERCHLORATE ION CONCENTRATION IN SODIUM HYPOCHLORITE SOLUTIONS AND

DRINKING WATER: KINETICS OF PERCHLORATE ION FORMATION AND EFFECTS OF ASSOCIATED CONTAMINANTS

by Aleksey N. Pisarenko

The dissertation consists of six chapters that summarize the investigation of

factors impacting perchlorate ion formation in sodium hypochlorite solutions and the

development of a predictive model for perchlorate ion formation. There are also two

appendices detailing the synthesis and applications of nanomaterials for designing

sensors.

Chapter 1 gives a brief history of perchlorate ion as an emerging contaminant. A

background on the occurrence, toxicology, and regulatory actions is provided.

Chapter 2 focuses on the analytical methods that were developed and validated for

analysis of perchlorate, bromate, chlorate, hypochlorite, and chlorite ions in various

hypochlorite ion solutions. Comparison of a LC-MS/MS method and an iodometric

titration method is provided. The chapter also details sample preparation methods, such

as the use of malonic acid to stop formation of perchlorate ion.

Chapter 3 details the experimental matrix to identify factors that impact

perchlorate ion formation. Effects of different contaminants were investigated at elevated

temperatures.

Chapter 4 provides a detailed investigation of the effects of concentration of

hypochlorite and chlorate ions, ionic strength, and temperature. The order of the

perchlorate ion formation with respect to hypochlorite and chlorate ions was determined.

A thorough investigation of the various effects led to derivation of a simple expression

that relates the effects of ionic strength and temperature on the second-order rate

constant.

Chapter 5 focuses on validation and application of the developed predictive

expression on various hypochlorite ion solutions. Bleach 2001 Predictive Model was

used to predict the decomposition of hypochlorite ion, and the output was used together

with the predictive expression developed in this work to predict formation of perchlorate

ion in hypochlorite ion solutions. Potential formation of perchlorate ion in stored

hypochlorite ion solutions is discussed, and recommendations to minimize formation of

perchlorate ion are provided.

Chapter 6 summarizes the findings of this Dissertation and provides conclusions

in the context of perchlorate ion contamination of drinking water when hypochlorite ion

is used as a disinfectant.

ANALYTICAL MEASUREMENTS AND PREDICTIONS OF PERCHLORATE ION CONCENTRATION IN SODIUM HYPOCHLORITE SOLUTIONS AND

DRINKING WATER: KINETICS OF PERCHLORATE ION FORMATION AND EFFECTS OF ASSOCIATED CONTAMINANTS

A DISSERTATION

Submitted to the Faculty of

Miami University in partial

fulfillment of the requirements

for the degree of

Doctor of Philosophy

Department of Chemistry and Biochemistry

by

Aleksey N. Pisarenko

Miami University

Oxford, Ohio

2009

Dissertation Directors: Dr. Gilbert E. Pacey and Dr. Gilbert Gordon

©

Aleksey N. Pisarenko 2009

iii

Table of Contents

List of Tables vii

List of Figures x

Dedication xviii

Acknowledgements xix

1. Introduction 1

1.1 Perchlorate Ion: Introduction 1

1.2 Perchlorate Ion: Toxicity and Regulation 3

1.3 Hypochlorite Ion Solutions as Potential Source of Perchlorate Ion 4

1.4 Research Objectives 5

2. Analysis and Sample Preparation of Hypochlorite Ion Solutions: Analytical

Methods Summary 7

2.1 Introduction to the Analysis of Sodium Hypochlorite Solutions 7

2.1.2 Transition metal ions: Co2+, Cu2+, Fe3+, Mn2+, and Ni2+ 9

2.1.3 Specific Conductance, Ionic Strength, and pH Measurements 10

2.2 Results and discussion 10

2.2.1 The LC-MS/MS Analysis of Perchlorate, Bromate, and Chlorate

Ions 10

2.2.2 Validation of LC-MS/MS Method for the Analysis of

Hypochlorite Solutions 12

2.2.3 Iodometric Titrations: Analysis of Chlorite, Chlorate, and

Hypochlorite Ions 18

2.2.3.1 Adam-Gordon Method 19

2.2.4 Method Selection for the Measurement of Chlorate Ion 21

2.2.5 Selection of Quenching Agent 24

2.2.5.1 Safety, Ease of Handling, Transport, and Stability 26

2.2.5.2 Ability to Quench Hypochlorite Ion Reproducibly 28

2.2.5.3 Impact on the Analysis of Bromate, Chlorate, and

iv

Perchlorate Ion 29

2.2.5.3 Quenching Agent Selection Summary 32

2.3 Conclusions 34

3. Experimental Design: Identifying Factors Impacting the

Perchlorate Ion Formation in Hypochlorite Ion Solutions 36

3.1 Experimental Matrix and Chemicals 38

3.2 Effect of Hypochlorite Ion Concentration 39

3.3 Effect of Chlorate Ion Concentration 41

3.4 Effect of Transition Metal Ions (Co2+, Cu2+, Fe3+, Mn2+, and Ni2+) 43

3.5 Effect of Noble Metal Ions (Ag+, Au+, Ir3+, Pd2+, and Pt2+) 44

3.6 Effect of Chlorite Ion Concentration 46

3.6.1 Combined Effect of Transition Metal Ions, Chlorite and

Bromide Ions 50

3.7 Effect of Bromide Ion and Bromate Ion Concentration 52

3.8 Effect of Ionic Strength 54

3.9 Effect of pH 57

3.10 Conclusions 60

4. Kinetics of Perchlorate Ion Formation and Determination

of the Rate Law 63

4.1 Reaction Order with Respect to Chlorate Ion:

ln (d[ClO4-]/dt) vs. ln [ClO3

-] 64

4.2 Reaction Order with Respect to Chlorate Ion:

ln (d[ClO4-]/dt) vs. ln [OCl-] 69

4.3 Multiple Reaction Pathways 75

4.3.1 Parallel Reaction Pathway 76

4.3.2 Consecutive Reaction Pathway 78

4.4 Ionic Strength Effect on the Rate of Perchlorate Ion Formation 80

4.4.1 Dependence of the Second-Order Rate Constant on

the Ionic Strength 84

v

4.4.2 Dependence of the Second-Order Rate Constant on

the Temperature 88

4.4.3 Combining the Effects of the Ionic Strength and Temperature

on the Second-Order Rate Constant 91

4.5 Conclusions 92

5. The Perchlorate Ion Formation Model: Validation and

Applications 93

5.1 Predicted Perchlorate Ion Formation in Bulk Sodium


5.2 Predicted Perchlorate Ion Formation in Real-World Bulk

Sodium Hypochlorite Solutions 98

5.3 Using Perchlorate Ion Formation Model to Determine Implications of Bulk Sodium Hypochlorite Solutions Storage 102 5.4 Application of the Perchlorate Model to OSG Sodium


5.5 Application of The Perchlorate Model to Calcium Hypochlorite

Solutions 112

5.6 Potential Contribution of Perchlorate Ion to Drinking Water 114

from Various Hypochlorite Ion Solutions 112

5.7 Conclusions 118

6. Conclusions 119

6.1 Summary 119

6.2 Recommendation to Water Utilities 120

Appendix 1. Detection of Ozone Gas by Gold Nanoislands 122

A1.1 Introduction 122

A1.2 Experimental 123

A1.3 Results and Discussion 124

A1.4. Conclusions 131

vi

Appendix 2. Electrochemically Assisted Processing of Organically

Modified, Perpendicularly Oriented Mesoporous Silica

Films with Fluorescent Functionality 132

A2.1. Introduction 132

A2-2. Experimental Details 136

A2-3. Results and Discussion 137

A2-4. Conclusions 142

A2.5 Acknowledgements 143

References 144

vii

List of Tables

Table 1. ICP-MS MRLs (µg/L) in water and hypochlorite ion solutions 9 Table 2. MDL data for perchlorate, bromate, and chlorate ions (n = 8) 13 Table 3. Spike recoveries of analytes with and without filtration and at different 15 Table 4. Standardization of sulfite and thiosulfate ion solutions by 0.109 M IO3

- 20 Table 5. Comparison of measurements by the LC-MS/MS and iodometric titration for bulk hypochlorite ion solutions (n = 7) 21 Table 6. Comparison of measurements by the LC-MS/MS and iodometric titration for OSG sodium hypochlorite solutions at less than 1.0 g/L ClO3

- (< 10 mM) (n ≥ 3) 22 Table 7. Effects of malonic acid (MA) on recoveries of chlorate, perchlorate, and bromate ions measured by LC-MS/MS (n = 3, replicate Samples Analyzed in triplicate; S.D. = standard deviation) 30 Table 8. Effects of malonic acid (MA) on analysis of perchlorate and bromate ions at different dilutions (n = 3; S.D. = standard deviation) 30 Table 9. Effects of quenching agent on analysis of chlorate comparison of LC-MS/MS and titration results (n = 3, replicate samples analyzed in triplicate; S.D. = standard deviation) 31 Table 10. Perchlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2-month period 31 Table 11. Bromate ion stability in hypochlorite ion solutions quenched with malonic acid over 2-month period 32 Table 12. Chlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2-month period 32 Table 13. Summary of quenching agent test results and decision-making matrix 33 Table 14. Changes in perchlorate ion concentration of samples spiked with Ag+, Au+, Ir3+

, Pd2+, and Pt2+ (Noble Me) vs. control (no spike), incubated at 50 ºC 46 Table 15. Decomposition of hypochlorite ion at 30 ºC in solutions, at various initial concentrations of chlorate ion ([ClO3

-]0) at pH ~12.5 66 Table 16. Reaction order with respect to chlorate ion and corresponding

viii

correlation coefficients in solutions at constant hypochlorite ion at pH ~ 12.5 and various temperatures 69 Table 17. Decomposition of hypochlorite ion at 30 ºC in solutions at constant chlorate ion concentration at pH ~12.5 71 Table 18. Parallel reaction pathway experimental rate constants in solutions, at various hypochlorite ion and constant chlorate ion at pH ~12.5 78 Table 19. Consecutive reaction pathway experimental rate constants at various initial concentrations of hypochlorite ion and constant chlorate ion at pH ~12.5 80 Table 20. Ionic strength (μ) of hypochlorite ion solutions at various chlorate ion at 40 ºC experiments (TDS = Total Dissolved Solids) 81 Table 21. Ionic strength (μ) of hypochlorite ion solutions at various hypochlorite ion at 40 ºC experiments (TDS=Total Dissolved Solids) 81 Table 22. Slopes and intercepts of least-squares lines shown in Figure 52 88 Table 23. Experimental and predicted second-order rate constants at variable ionic strength and temperature (kexp = experimental k2; kpred = predicted k2) 91 Table 24. Predicted changes in hypochlorite ion, chlorate ion, and d[ClO4

-]/dt as a function of time at 40 ºC 96 Table 25. Bromate and perchlorate ions, and transition metals in bulk utility hypochlorite ion solutions 98 Table 26. Chlorate and hypochlorite ions, pH, TDS, and ionic strength in bulk utility hypochlorite ion solutions 99 Table 27. Transition metals, bromate, chlorate, and perchlorate ions in OSG

hypochlorite ion solutions 108 Table 28. Concentration of hypochlorite ion, pH, TDS, and ionic strength (μ) in OSG hypochlorite ion solutions 109 Table 29. Second-order rate constants of perchlorate ion formation in OSG hypochlorite ion solutions, experiment vs. model (k2obs= experimental k2, k2cal = predicted k2, in units of L·mol-1·day-1) 111 Table 30. Second-order rate constants of perchlorate ion formation in calcium hypochlorite solutions, experiment vs. model (k2obs= experiment k2, k2cal = k2 predicted in units of L·mol-1·day-1) 114 Table 31. Residence time of the sampled distribution waters 115 Table 32. Perchlorate ion in raw, finished, and distribution waters 115 Table 33. Bromate ion in raw, finished, and distribution waters 115 Table 34. Chlorate ion in raw, finished, and distribution waters 115

ix

Table 35. Contributions of perchlorate, bromate, and chlorate ions per mg FAC in various hypochlorite ion solutions 117

x

List of Figures Figure 1. Perchlorate ion 1

Figure 2. Extracted ion chromatogram of perchlorate ion MRM m/z 98.9/82.8

of standard solution containing 0.02 μg/L of perchlorate ion 13


of (a) sodium hypochlorite solution with 0.1 μg/L ClO4-; (b) standard

solution with 0.1 μg/L ClO4- 14

Figure 4. Bromate ion chromatograms of sodium hypochlorite sample diluted

by: (a) factor of 1:10; (b) factor of 1:100; (c) factor of 1:1000 16

Figure 5. Actual sample concentrations of analytes measured at different

dilutions 17

Figure 6. Comparison of chlorate ion measurements by (a) iodometric titration

and (b) by the LC-MS/MS, during a chlorate ion spike experiment,

at 75 ºC (Control = solution at initial [OCl-] = 1.46 M and [ClO3-]

= 0.29 M) 23

Figure 7. Stock solutions of ascorbic acid freshly prepared (left), after 20 days

(center), and after 37 days of storage (right) 27

Figure 8. Ascorbic acid -quenched hypochlorite ion sample solutions (left 3

bottles) and malonic acid-quenched solution (right bottle) 28

Figure 9. Chromatogram of bromate (left) and 18O-labeled bromate (right) of

(a) sulfite-quenched sample, and (b) thiosulfate-quenched sample of

13% sodium hypochlorite solution diluted by a factor of 1:10,000 29

Figure 10. Decomposition of hypochlorite ion and formation of chlorate ion at

75 ºC in solutions, at various initial concentrations of hypochlorite

ion 40

Figure 11. Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions,

at various initial concentrations of hypochlorite ion 40


75 ºC in solutions, at various initial concentrations of chlorate ion 42

Figure 13. Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions,

at various initial concentrations of chlorate ion 42

xi

Figure 14. Effects of Transition Metals Ions (Me = Co2+, Cu2+, Fe3+, Mn2+,

and Ni2+) on the hypochlorite ion decomposition and the chlorate

ion formation 43

Figure 15. Effects of Transition Metals Ions (Me = Co2+, Cu2+, Fe3+, Mn2+,

and Ni2+) on perchlorate ion formation 44

Figure 16. Effects of noble metal ions (Noble Me = Ag+, Au+, Ir3+, Pd2+,

and Pt2+), plots of hypochlorite ion decomposition and chlorite ion

formation 45

Figure 17. Effects of noble metals ions (Noble Me = Ag+, Au+, Ir3+, Pd2+,

and Pt2+), overlaid plots of perchlorate ion formation 45

Figure 18. Plots of hypochlorite ion decomposition and chlorate ion formation, in

solutions at various initial concentrations of chlorite ion and/or chlorate

ion at (a) 50 ºC, (b) 30 ºC 47

Figure 19. Overlaid plot of changes in molar product and perchlorate ion

formation over time, in solutions at various initial concentrations of

chlorite ion and/or chlorate ion at (a) 50 ºC, (b) 30 ºC 48

Figure 20. Plots of perchlorate ion formation in solutions at various initial

concentrations of chlorite ion and/or chlorate ion at (a) 30 ºC,

(b) 50 ºC 49


50 ºC in solutions spiked with bromide, chlorite, and transition metal

ions (Me = Co2+, Cu2+, Fe3+, Mn2+, and Ni2+) 51

Figure 22. Formation of perchlorate ion at 50 ºC in solutions spiked with bromide,

chlorite, and transition metal ions (Me = Co2+, Cu2+, Fe3+, Mn2+,

and Ni2+) 51

Figure 23. Decomposition of hypochlorite ion at 50 ºC in solutions spiked with

bromide and bromate ions at pH~12.5 53

Figure 24. Formation of bromate ion at 50 ºC in solutions spiked with bromide

and bromate ions at pH~12.5 53


and bromate ions at pH~12.5 54

xii

Figure 26. Decomposition of hypochlorite ion at 40 ºC in solutions at various

initial concentrations of chloride ion 56

Figure 27. Formation of chlorate ion at 40 ºC in solutions at various initial

concentrations of chloride ion 57

Figure 28. Formation of perchlorate ion at 40 ºC in solutions at various initial

concentrations of chloride ion 57


40 ºC in (a) 1.4 M OCl-, (b) 0.9 M OCl- solutions at various initial

pH 59

Figure 30. Formation of perchlorate ion at 40 ºC in (a) 1.4 M OCl-, (b) 0.9 M

OCl- solutions at various initial pH 60

Figure 31. Overlaid plots of (a) hypochlorite ion decomposition; (b) chlorate

ion formation at 60 ºC 0.12 M OCl- solutions at various

initial pH 61

Figure 32. Overlaid plots of perchlorate ion formation at 60 ºC 0.12 M OCl-

solutions with various initial pH 62

Figure 33. Overlaid plots of (a) perchlorate ion formation; (b) decomposition

Of hypochlorite ion and formation of chlorate ion at 30 ºC in

solutions at various initial concentrations of chlorate ion at pH ~12.5 65


of hypochlorite ion and formation of chlorate ion at 40 ºC in





Figure 36. Fitted natural log lines of the rate of perchlorate ion formation as

a function of chlorate ion concentration at 30 ºC in solutions

constant in hypochlorite ion at pH ~12.5 67




xiii






solutions at various initial concentrations of hypochlorite ion

at pH ~12.5 70




at pH ~12.5 70




at pH ~12.5 71


a function of hypochlorite ion concentration at 30 ºC in solutions

at constant chlorate ion at pH ~12.5 72












solutions at constant molar product at pH ~12.5 74


xiv


solutions at constant molar product at pH ~12.5 75

Figure 48. Parallel reaction pathway linear fitted plots of (a) 30 ºC

experiment; (b) 40 ºC experiment; (c) 50 ºC experiment; in

solutions at constant chlorate ion and various hypochlorite ion

at pH ~12.5 77

Figure 49. Consecutive reaction pathway linear fitted plots of (a) 30 ºC

experiment; (b) 40 ºC experiment; (c) 50 ºC experiment; in

solutions at constant chlorate ion and various hypochlorite ion

at pH ~12.5 79

Figure 50. Rate of perchlorate ion formation as a function of (a) ionic

strength; (b) concentration of hypochlorite ion at 40 ºC 82

Figure 51. Plot of )1/( μμ + and μb term as a function of ionic

strength. Note: value of 0.5 was assumed for the b term as an

approximation 86

Figure 52. Overlaid linear plots of log of second-order rate constant versus

ionic strength in solutions at various initial concentrations of

hypochlorite ion at different temperatures 87

Figure 53. Smooth-line plots of the perchlorate ion formation as a function

of time in solutions at similar initial hypochlorite and chlorate ions

and various temperatures 88

Figure 54. Linear plot of ln(k0/T) as a function of (1/T) 90

Figure 55. Smoothed-line plots of hypochlorite ion decomposition and

chlorate ion formation determined experimentally in conjunction

with Bleach 2001 (Error bars set at ± 10 %), for solutions at (a)

[OCl-]0 = 82 g/L, [ClO3-]0 = 63 g/L at 30 ºC; (b) [OCl-]0 = 70 g/L,

[ClO3-]0 = 51 g/L at 40 ºC; (c) [OCl-]0 = 83 g/L, [ClO3

-]0 = 50 g/L

at 50 ºC 94

Figure 56. Overlaid smoothed-line plots of predicted (Error bars set at ± 10

%) perchlorate ion formation and determined experimentally,

for solutions incubated at (a) 30 ºC; (b) 40 ºC; (c) 50 ºC 97

xv

Figure 57. Overlaid smoothed-line plots of (a) hypochlorite ion

decomposition; (b) chlorate ion formation; (c) perchlorate ion

formation determined experimentally during incubation at 50 ºC 99

Figure 58. Overlaid smoothed-line plots of hypochlorite ion decomposition

and chlorate ion formation determined experimentally in

conjunction with Bleach 2001 (Error bars set at ± 10%), for

solutions with (a) [OCl-]0 = 63 g/L, [ClO3-]0 = 23 g/L; (b) [OCl-]0 =

111 g/L, [ClO3-]0 = 8.7 g/L; (c) [OCl-]0 = 89 g/L, [ClO3

-]0 = 4.4 g/L ;

(d) [OCl-]0 = 97 g/L, [ClO3-]0 = 12 g/L; incubated at 50 ºC 100

Figure 59. Overlaid smoothed-line plots of predicted (Error bars set at ± 10

%) perchlorate ion formation and determined experimentally at

50 ºC in solutions with (a) [OCl-]0 = 63 g/L, [ClO3-]0 = 23 g/L;

(b) [OCl-]0 = 111 g/L, [ClO3-]0 = 8.7 g/L; (c) [OCl-]0 = 89 g/L,

[ClO3-]0 = 4.4 g/L; (d) [OCl-]0 = 97 g/L, [ClO3

-]0 = 12 g/L 101

Figure 60. Smoothed-line plot of the rate of perchlorate ion formation as a

function of (a) temperature; (b) dilution factor. Note: Rate in

solution at 2.54 M OCl-, 0.034 M ClO3-, and μ = 7.5 M; rate at

35 ºC = 100% 103

Figure 61. Overlaid smoothed-line plots of (a) predicted decomposition of

hypochlorite ion and formation of perchlorate ion; (b) plot of μg

ClO4- per mg OCl- as a function of time in solutions at 2.03

M OCl- and 1.02 M OCl- at 35 ºC 105


hypochlorite ion and formation of perchlorate ion; (b) plot of μg

ClO4- per mg OCl- as a function of time in solutions at 2.03

M OCl- and 1.02 M OCl- at 25 ºC 106

Figure 63. Overlaid smoothed-line plots of hypochlorite ion decomposition

(a) OSG solutions 1-6; (b) OSG solutions 7-12, 50 ºC 109

Figure 64. Overlaid smoothed-line plots of perchlorate ion formation (a)

OSG solutions 1-6; (b) OSG solutions 7-12, 50 ºC 110

Figure 65. Smoothed-line plots of hypochlorite ion decomposition and

xvi

chlorate ion formation in calcium hypochlorite solutions (a)

incubated at 50 ºC; (b) incubated at 60 ºC 113

Figure 66. Smoothed-line plots of perchlorate ion formation in calcium

hypochlorite solutions (a) incubated at 50 ºC; (b) incubated at

60 ºC 113

Figure A1-1. SEM images of typical gold nanoislands produced by sputtering

process on a polished aluminum substrate 124

Figure A1-2. AFM images of 25 nm (a); and 14 nm (b) gold nanoislands on

quartz substrate 125

Figure A1-3. Overlaid UV-Vis spectra of : dashed line-gold thin film on

quarts, solid line-gold nanoislands, dot-dashed line gold

nanoislands exposed to ozone gas, dotted line-gold nanoislands

reversed by annealing at 375 ºC for 15 min 126

Figure A1-4. Overlaid UV-Vis spectra of 25 nm gold nanoislands with surface

Plasmon absorbance max at 520 nm exposed to concentrations

of ozone, increased in increments form 20.9 μg/L to 166.1μg/L.

Ozone causes a red-shift in the surface-plasmon absorbance max 128

Figure A1-5. Shift of the 25 nm gold nanoislands surface-plasmon max (520

nm) as a function of ozone concentration, logarithmic fit gives

an equation of y=6.8ln(x)-15.19 produced a correlation

coefficient of 0.9659 129

Figure A1-6. Surface plasmon’s shifts of gold nanoislands with absorbance max

At 532 nm as a function of ozone concentration 130

Figure A2-1. (a) Left sample—Blank ITO Electrode, right sample—ITO

Electrode with the EPON film. (b) Samples in same position

under UV light, fluorescence is observed for ITO Electrode with

EPON film 137

Figure A2-2. EPON-Coated ITO Electrode, plating time 10 s at -2.1V. Imaging

of the plating interface shows the difference in surface

morphology that of the ITO and that of EPON film, which

indicates EPON film is deposited on the ITO surface 138

xvii

Figure A2-3. EPON-coated ITO electrode, plating time: (a) 30 s -2.1V. The

Deposited film indicates normal to the electrode surface orientation

of the deposited EPON film. (b) EPON-coated ITO electrode,

magnification of (a) reveals normal orientation of the mesopores 139

Figure A2-4. (a) Dry Sol-Gel Film; (b) Dry, after CTAB and free polymer

extraction; (c) Wetted with water; (d) Dry. (b) and (d) overlap. The

difference in emission intensity between (a) and (b) amounts to

removed 4-methylcoumarin-7-yl 3-(trimethoxysilyl)

propylcarbamate not bound to EPON film. (b) and (d) are the same

film before and after wetting, where (c) shows intensity drop when

the obtained film is washed of CTAB and 4-methylcoumarin-7-yl

3-(trimethoxysilyl)propylcarbamate, that is not bound; dried and

wetted again 140

Figure A2-5. Overlaid excitation spectra of EPON film subjected to different

pH: (a) DI water; (b) pH 1; (c) pH 2.2; (d) pH 13.3. Excitation

maximum and excitation peak shape shifts based on pH of the

wetting solution 141

Figure A2-6. Overlaid emission spectra of EPON film subjected to different

pH: (a) DI water; (b) pH 1; (c) pH 2.2; (d) pH 13.3 Emission

maximum and peak shape, consistent with excitation peak

changes, shifts based on pH of the wetting solution 142

xviii

Dedication

For grandmothers Tamara Vaganova and Nina Pisarenko

Both regrettably ahead of their time

For my dear parents, Nikolai and Liubov Pisarenko,

my sister, Liubov Pisarenko

And last but not least, For Daniel S. Elliott

xix

Acknowledgements

I would like to thank Miami University’s Chemistry and Biochemistry

Department for the continuous support of my graduate studies. I also wish to thank

Southern Nevada Water Authority for selecting me as a graduate intern and allowing me

to be a part of a great research team at Applied Research and Development Center. In

addition, I am grateful to the funding made available through American Water Works

Association (AWWA), Water Research Foundation (WRF), and The Ohio Third Frontier

IDCAST Wright Center for Innovation. I would like to give special thanks to members

of my committee:

To Dr. Gilbert Pacey and Dr. Gilbert Gordon, my Dissertation Directors, who

helped me to shape my graduate career by providing me with freedom to explore new

topics, motivation, and sometimes pressure. This invaluable experience I will use for the

rest of my life. I would like to thank Dr. Pacey for letting me to take on a number of

opportunities and various research projects that have been both very educational and

rewarding in the end. I am very grateful to Dr. Gilbert Gordon for entrusting me to lead

on numerous projects which have led me to great opportunities and have enormously

broadened my understanding of chemistry. Thank you both for enhancing my graduate

career and for being my mentors.

To Dr. Luis Actis, and Dr. Richard Taylor, who have provided contributions to

promote both my research and academic progress, by research collaborations and useful

suggestions to help me stay on-track.

To Dr. Michael Crowder, for your willingness to be part of my committee at later

stages of my research project. I also would like to specially thank you for helping me to

select Miami University and for your help with my transition.

Also, I would like to acknowledge several people that have helped me along the

way during the past several years. I would like to thank Dr. Wolfgang Spendel for his

creative ideas and discussions that have significantly contributed to the work described in

the appendices. I would like to thank several faculty members and their group members:

Dr. James Cox for his useful discussions and help with the project involving sol-

gel and electrochemical methods (Appendix 1). I would like to thank Dr. Diep Ca for

xx

help with ICP-OES training and Kamila Wiaderek, with whom I often shared chemicals

and equipment. Dr. Richard Taylor and Jordan Brown, an undergraduate student

working in Dr. Taylor’s lab, both have significantly contributed to the work described in

Appendix 2. I thank Dr. Shouzong Zou for sharing chemicals and equipment and Dr.

Sachin Kumar for useful discussions on the topic of metal nanoparticles and the use of

AFM. Also I wish to thank Dr. Thomas Riechel and Dr. Neil Danielson for their help and

guidance with the teaching assignments and for their help with instrument

troubleshooting.

In addition I would like to thank several staff members of Chemistry and

Biochemistry department. I wish to thank Dr. Ian Peat for his help with running /

troubleshooting ICP-MS, and for the training to perform analysis by MALDI-TOF and

LC-MS/MS. I wish to acknowledge Barry Landrum for the undisputedly-superb

machinist skills to craft specialty-designed cell holders, adaptors, and other nifty devices,

that have significantly enhanced the lab work and reduced costs. Lynn Johnson has been

very helpful with troubleshooting a number of electrical problems and I would like to

thank him for his efforts and time. I also would like to thank Dr. Hans Bier for

interesting discussions and for his help with finding the “right equipment for the job.” I

also thank Dianna (Deedee) Bear for her logistical help with the teaching assignments,

and Kim Traylor for her help with purchasing chemicals and equipment. I also wish to

acknowledge Shelli Minton and Sharon Weber for being very helpful finding solutions to

problems one encounters outside the research lab.

I wish to acknowledge Richard Edelman and Mathew Dewly for their help with

the training on SEM and TEM, and for their useful suggestions to enhance the analysis.

I also would like to mention several friends that I have acquired while attending

Miami University. Dr. Justin Heuser and Dr. Sean Pucket, who consistently challenged

my “background” but otherwise, have contributed to my general progress through

graduate school. I would like to recognize Dr. Heuser for putting up with me as a

roommate and for being a good friend. It was also my pleasure to have met Dr. Anita

Taulbee-Combs, Kamila Wiaderek, Olaf Borkiewicz, Dr. Sachin Kumar, Dr. Peter Xu,

Dr. Pattraranee Limphong, Patrick Hensley, Sriram Devanathan, Josh Ebel, Jordan

Brown, Matt Bachus, David Hufnagle, and many others.

xxi

I am also very grateful to all of the people at SNWA Applied Research and

Development Center that I had a great pleasure working with and with whom I have

learnt so much. I would like to specially thank Dr. Benjamin Stanford, Dr. Shane Snyder,

and Dr. Gilbert Gordon, for their guidance and experience that has helped me to stay on

track in many ways. In addition I would like to thank Oscar Quiñones, Dr. Douglas

Mawhinney, Brett Vanderford, and Rebecca Trenholm for their assistance with the

development and troubleshooting of the analytical methods employed at SNWA. I am

also thankful to Janie Holady, Shannon Fergusson, Elaine Go, and Christy Meza for the

assistance with the sample handling and preparation. I also wish to thank Dave Rexing

and Linda Parker for their assistance and support of the project.

Lastly, I would like to thank an alumnus of Miami University, Dr. Luke Adam,

for his extensive work on hypochlorite ion solutions with Dr. Gilbert Gordon, which has

greatly enhanced the investigation of perchlorate ion formation.

Adapted from "Hypochlorite—An Assessment of Factors That Influence the

Formation of Perchlorate and Other Contaminants," by permission. Copyright ©

2009, American Water Works Association, co-sponsored by AWWA Water

Industry Technical Action Fund (WITAF), the Water Research Foundation, and

the Southern Nevada Water Authority.

1

Figure 1. Perchlorate ion

CHAPTER 1. INTRODUCTION

1.1 Perchlorate Ion: Introduction

In 1997, the California Department of Health Services (now California

Department of Public Health) began sampling for perchlorate ion, prompting the

Sanitation and Radiation Laboratory to develop a sensitive analytical method for

measurement of perchlorate ion in water, achieving a detection limit of 4 μg/L. 1 At the

time, the method constituted a major improvement in sensitivity, allowing detection of

perchlorate ion in drinking water wells. This finding resulted in a most rigorous review

of the issue of perchlorate ion contamination by the scientific community and the

regulatory agencies to this day.2

Perchlorate ion consists of a tetrahedral array of

oxygen atoms with a central chlorine atom, structure

shown in Figure 1. The chlorine atom is at it highest

“formal” oxidation state of “+7,” thus the species are a

strong oxidizing agent. However, because the

reduction-oxidation behavior of the perchlorate ion is

rarely observed and because perchlorate ion has less

tendency to form metal complexes than other anions, it has been commonly used as an

electrolyte to probe ionic equilibria in aqueous solutions.3 Reduction of perchlorate ion,

shown by Equation 1, is very slow and is typically observed only in presence of

concentrated acid.4

ClO4- + 8H+ + 8e- = Cl- + 4H2O Eo = 1.287 V at pH 0 (1)

1 CDHS, "Determination of Perchlorate by Ion Chromatography", California Department of Health Services (CDHS), 1997, Retrieved from: http://www.cdph.ca.gov/certlic/drinkingwater/Documents/Perchlorate/SRLperchloratemethod1997.pdf 2 CDPH, "History of Perchlorate in California Drinking Water", California Department of Public Health (CDPH), 2007, Retrieved 07/28/09, from: http://www.cdph.ca.gov/certlic/drinkingwater/Pages/Perchloratehistory.aspx 3 Robinson, R. A. and Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths Publications Limited: London, 1959. 4 Housecroft, C. E. and Sharpe, A. G. Inorganic Chemistry; Pearson Education Limited, 2001.

2

The occurrence of perchlorate ion in the environment has both natural and

anthropogenic origins. Fertilizers have been suspected to contain perchlorate ion.5 The

nitrate rich mineral, known as Chilean Saltpeter, which was found to contain trace

amounts of perchlorate ion,6 is used as a fertilizer in the United States and may have

introduced perchlorate ion contamination into soil and water. Other fertilizer materials

such as phosphate rock, potash, and ammonium dihydrogen phosphate have been also

reported to contain perchlorate ion,7 although the concentration of the perchlorate ion in

these fertilizers has not been attributed to a significant source of environmental

contamination.8 Recently it has been found that perchlorate ion may form naturally by

photochemical oxidation of chloride ion under ultraviolet (UV) light9 and in the presence

of atmospheric ozone.10

Perchlorate ion can be formed by electrolytic production,11 and perchlorate salts

can be isolated; however, they are very explosive and require special care. For example

ammonium perchlorate has been used in pyrotechnics, ammunition, and as a solid rocket

fuel.12 About 825 tons of ammonium perchlorate are used by the space shuttle booster

rockets. Production of perchlorate salts ramped from 2,000 tons in the 1950s to the peak

of 15,000 ton in mid-1980s.13 There has been at least one accident associated with the

production of perchlorate. In 1988, a large ammonium perchlorate plant, owned by this

Pacific Engineering & Production Co. of Nevada (PEPCON), exploded, and the

explosion is suspected of causing a release of perchlorate ion into the environment.

Ten years later, perchlorate ion was detected in the Colorado river in California,

traced to Lake Mead, and eventually to sites associated with ammonium perchlorate

manufacturing. In 2005, USEPA updated a map of perchlorate ion occurrence 5 Renner, R. "Study finding perchlorate in fertilizer rattles industry" Environ. Sci. Technol 1999, 33, 394A-395A. 6 Urbansky, E. T., Brown, S. K., Magnuson, M. L. and Kelty, C. A. "Perchlorate levels in samples of sodium nitrate fertilizer derived from Chilean caliche" Environ. Pollut. 2001, 112, 299-302. 7 Urbansky, E. T., Collette, T. W., Robarge, W. P., Hall, W. L., Skillen, J. M. and Kane, P. F. Survey of fertilizers and related materials for perchlorate (Cl04

-), EPA, 2001. 8 Urbansky, E. T. "Perchlorate as an environmental contaminant" Env. Sci. Pollut. Res. 2002, 9, 187-192. 9 Kang, N., Anderson, T. A. and Andrew Jackson, W. "Photochemical formation of perchlorate from aqueous oxychlorine anions" Anal. Chim. Act. 2006, 567, 48-56. 10 Kang, N., Jackson, W. A., Dasgupta, P. K. and Anderson, T. A. "Perchlorate production by ozone oxidation of chloride in aqueous and dry systems" Sci. Total Environ. 2008, 405, 301-309. 11 Mack, E. L. "Electrolytic Formation of Perchlorate" J. Phys. Chem. 1917, 21, 238-264. 12 Davis, C. O. Perchlorate explosives, E. I. du Pont de Nemours & Co., 1940, USA US 2190703 13 Motzer, W. "Perchlorate: Problems, Detection, and Solutions" Environ. Forensics 2001, 2, 301-311.

3

throughout the country.14 In California, these areas are associated with the facilities that

have manufactured or tested solid rocket fuels for the Department of Defense or the

National Aeronautics and Space Administration.15 In the 2005 Report on Perchlorate, the

United States Government Accountability Office (USGAO) reported the occurrence of

perchlorate ion in ground water, surface water, soil, and drinking water at 395 sites across

35 states.16 Once in water, perchlorate ion appears to be quite stable according to a

previous study,17 and thus is a long-term environmental concern.

1.2 Perchlorate Ion: Toxicity and Regulation

Perchlorate ion is an endocrine-disrupting compound, which inhibits the uptake of

iodine by the thyroid.18 The effects of perchlorate ion on the function of the thyroid and

hormone production in pregnant women are a critical issue to understand in order to

formulate a safe reference dose (RfD).19 In 2005 the National Academy of Sciences and

the United States Environmental Protection Agency (USEPA) established an RfD of

0.0007 mg/kg/day, which corresponds to a drinking water equivalent level (DWEL) of

24.5 µg/L.20 However, perchlorate ion is found not only in waters but also in food

sources.21 Thus, recognizing that there are multiple intake sources of perchlorate ion, the

USEPA issued in December, 2008 an interim health advisory level of 15 μg/L (based on

14 USEPA, "Known Perchlorate Releases in the U.S. - March 25th, 2005", United States Environmental Protection Agency, 2005, Retrieved 08/01/09, from: http://www.epa.gov/swerffrr/pdf/detect0305.pdf 15 Tikkanen, M. W. "Development of a drinking water regulation for perchlorate in California" Anal. Chim. Act. 2006, 567, 20-25. 16 USGAO Perchlorate: A System to Track Sampling and Cleanup Results Is Needed United States Government Accountability Office (USGAO): Washington, D.C., 2005. 17 Stetson, S. J., Wanty, R. B., Helsel, D. R., Kalkhoff, S. J. and Macalady, D. L. "Stability of low levels of perchlorate in drinking water and natural water samples" Anal. Chim. Act. 2006, 567, 108-113. 18 Greer, M. A., Goodman G. G., Pleus R. C, Greer S. E. "Health effects assessment for environmental perchlorate contamination: the dose response for inhibition of thyroidal radioiodine uptake in humans" Environ. Health Perspect. 2002, 110, 927-937. 19 Strawson, J., Zhao, Q. and Dourson, M. "Reference dose for perchlorate based on thyroid hormone change in pregnant women as the critical effect" Regul. Toxicol. Pharmacol. 2004, 39, 44-65. 20 USEPA Draft Regulatory Determinations Support Document for Selected Contaminants from the Second Drinking Water Contaminant Candidate List (CCL2) EPA Report 815-D-06-007, United States Environmental Protection Agency (USEPA): Washington, DC. , 2006. 21 USFDA, "2004-2005 Exploratory Survey Data on Perchlorate in Food", United States Food & Drug Administration, 2005, Retrieved 08/02/09, from: http://www.fda.gov/Food/FoodSafety/FoodContaminantsAdulteration/ChemicalContaminants/Perchlorate/ucm077685.htm

4

the RfD of 0.007 mg / kg but taking into account perchlorate ion intake from sources

other than water). However perchlorate ion is not regulated nationally in drinking water

at this time.22 Recently, perchlorate ion has been identified as a contaminant of concern

in sodium hypochlorite, widely used as a disinfectant for drinking water.23,24 Given a

growing concern of perchlorate ion exposure from multiple sources, it is not clear

whether the USEPA will reconsider regulation of perchlorate ion in drinking water.25

1.3 Hypochlorite Ion Solutions as Potential Source of Perchlorate Ion

In light of anticipated heightened security requirements by Department of

Homeland Security for the use of chlorine gas, water utilities may consider the use of

sodium hypochlorite as an alternative, to disinfect and maintain a chlorine residual level

of disinfectant in drinking water and waste water treatment applications.26 According to

a 2007 survey, 63 % of water treatment facilities use chlorine as the primary disinfectant.

As many as 30% of the drinking water treatment facilities (DWTF) in North America had

switched from using chlorine gas to hypochlorite ion solutions in the past 10 years.27

Chlorine gas was replaced in 81 % of the treatment facilities by bulk sodium hypochlorite

solutions, 17 % of sites switched to the use of on-site generated (OSG) sodium

hypochlorite solutions, and around 1 % of DWTF use calcium hypochlorite.

In addition to the currently unregulated chlorate ion, several regulated

contaminants are present in hypochlorite ion solutions including bromate ion,28,29

22 USEPA, "Interim Drinking Water Health Advisory for Perchlorate", United States Environmental Protection Agency, 2008, Retrieved 07/27/09, from: http://www.epa.gov/OGWDW/contaminants/unregulated/pdfs/healthadvisory_perchlorate_interim.pdf 23 Greiner, P., Mclellan, C., Bennet, D. and Ewing, A. "Occurrence of perchlorate in sodium hypochlorite" J. Am. Water Work Assoc. 2008, 100, 68-74. 24 Asami, M., Kosaka, K. and Kunikane, S. "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply" J. Water Supply Res. Technol. 2009, 58, 107-115. 25 Renner, R. "EPA perchlorate decision flawed, say advisers" Environ. Sci. Technol. 2009, 43, 553-554. 26 Shah, J. and Qureshi, N. "Chlorine Gas vs. Sodium Hypochlorite: What's the Best Option" Opflow 2008, 24-27. 27 Routt, J., Mackey, E. and Noack, R. "Committee Report: Disinfection Survey, Part 2 - Alternatives, experiences, and future plans" J. Am. Water Work Assoc. 2008, 100, 110-124. 28 Asami, Kosaka and Kunikane "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply". 29 Chlorine Institute, "Bromate in Sodium Hypochlorite Potable Water Treatment," 2004, Retrieved 04/14/09, from: http://www.chlorineinstitute.org/files/FileDownloads/BromateinNaOCl-PotableWaterTreatment.pdf

5

hypochlorite ion itself, and chlorite ion.30 Occurrence of bromate ion in drinking water

has also been linked to the chlorination process.31,32

Perchlorate ion, recently identified as a contaminant, has been shown to increase

during storage in sodium hypochlorite solutions. 33,34 Interestingly, high concentrations

of perchlorate ion in stored sodium hypochlorite solutions were also correlated to high

concentrations of chlorate ion,35 though the factors impacting the rate of formation were

not well understood at the time.

Because sodium hypochlorite is already widely used in drinking water treatment

and the growing number of water treatment facilities that have switched to hypochlorite

ion solutions36 for chlorination and chloramination, the contribution of perchlorate ion

from hypochlorite ion solutions is a critical issue.

1.4 Research Objectives

Sodium hypochlorite is generally manufactured by passing chlorine gas through

sodium hydroxide, which is typically produced by means of the chlor-alkali process, in

which an aqueous sodium chloride solution is electrolyzed to produce chlorine gas and

sodium hydroxide.37 The decomposition of the bulk alkaline hypochlorite ion solutions

has been extensively studied,38,39 and as a result, a predictive hypochlorite ion

30 Gordon, G., Adam, L. C., Bubnis, B., Hoyt, B., Gillette, S. J. and Wilczak, A. "Controlling the formation of chlorate ion in liquid hypochlorite feedstocks" J. Am. Water Work Assoc. 1993, 85, 89-97. 31 Weinberg, H. S., Delcomyn, C. A. and Unnam, V. "Bromate in Chlorinated Drinking Waters: Occurrence and Implications for Future Regulation" Environ. Sci. Technol. 2003, 37, 3104-3110. 32 Bouland, S., Duguet, J. P. and Montiel, A. "Evaluation of bromate ions level introduced by sodium hypochlorite during postdisinfection of drinking water" Environ. Technol. 2005, 26, 121-125. 33 Greiner, Mclellan, Bennet and Ewing "Occurrence of perchlorate in sodium hypochlorite". 34 Asami, Kosaka and Kunikane "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply". 35 Asami "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply". 36 Routt, Mackey and Noack "Committee Report: Disinfection Survey, Part 2 - Alternatives, experiences, and future plans". 37 Chlorine Institute, "Pamthlet 96 Sodium Hypochlorite Manual, Edition 3," 2006, Retrieved 07/08/09, from: http://www.chlorineinstitute.org/files/Pamphlet096Edition3April2006FinalWebsite.pdf 38 Adam, L. C. An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite Ph.D. Dissertation, Miami University, Oxford, OH, 1994. 39 Adam, L. C. and Gordon, G. "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion" Inorg. Chem. 1999, 38, 1299-1304.

6

decomposition model was developed (referred as Bleach 2001 in later chapters).40 As a

result of these studies, it was found that hypochlorite ion solutions are most stable in the

pH range of 11-13 and at lower temperatures. Decomposition of hypochlorite ion, in

alkaline solutions, is second-order resulting in production of oxygen gas, chlorate, and

chloride ions.41 The presence of transition metals such as Ni(II) and Cu(II) were found to

catalyze decomposition of hypochlorite ion.42

Thus a series of objectives were proposed to investigate factors impacting the

formation of perchlorate ion in hypochlorite ion solutions:

i. Determine which analytical methods are most suitable for analysis of

concentrated sodium hypochlorite solutions to determine concentration of hypochlorite,

chlorite, chlorate, perchlorate, and bromate ions. Because perchlorate ion has been

reported in hypochlorite ion solutions in concentrations below 1 mg/L, a sensitive and

accurate method is needed.

ii. To stop formation of the perchlorate ion in hypochlorite ion solutions, a

quenching agent needs to be selected on the basis of minimum effects on the analytical

measurements.

iii. Determine whether the transition metal ions, noble metal ions, bromate

and chlorite ions catalyze formation of perchlorate ion.

iv. Determine concentration effects of chlorate and hypochlorite ions on the

rate of formation of perchlorate ion.

v. Determine the reaction order of perchlorate ion formation and propose the

simplest rate law. Develop an easily followed model that can predict the formation of the

perchlorate ion in the concentrated sodium hypochlorite solutions with an accuracy of ±

10 %.

vi. Determine strategies to minimize perchlorate ion formation in

concentrated sodium hypochlorite solutions. Develop recommendations to water utilities

that use hypochlorite ion for disinfection.

40 Adam, L., Gordon, G. and Pierce, D., Bleach 2001 Predictive Model, Miami University, Oxford, Ohio. Copyright (c) 2001 AwwaRF and AWWA. 41 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 42 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

7

CHAPTER 2. ANALYSIS AND SAMPLE PREPARATION OF

HYPOCHLORITE ION SOLUTIONS: ANALYTICAL

METHODS SUMMARY

The main objective in this portion of the study was to optimize and validate a

previously reported LC-MS/MS method for the simultaneous identification and

quantification of perchlorate, bromate, and chlorate ions in hypochlorite ion solutions. It

was hypothesized that a new, modified LC-MS/MS method would allow shorter run

times while offering higher sensitivity. Because this new method was to be used in the

subsequent investigation into the kinetics of perchlorate ion formation in hypochlorite ion

solutions, a robust, accurate method was required. Thus, an additional objective was to

compare LC-MS/MS methods to other validated methods, such as iodometric titration, to

identify concentration ranges at which the methods provided the most reliable results for

analysis of chlorate ion. Lastly, the selection of a hypochlorite ion quenching agent for

sample preservation was needed, and any effects on analysis were investigated.

2.1 Introduction to the Analysis of Sodium Hypochlorite Solutions

Mass spectrometry (MS) most commonly interfaced with ion-chromatography

(IC-MS)43, 44 or liquid chromatography (LC-MS)45,46 has been utilized for the detection

and quantification of oxyhalide anions. Both US EPA method 331.0, that uses LC with

tandem mass spectrometry (LC-MS/MS),47 and US EPA method 332.0, that uses IC-

43 Jackson, P. E., Gokhale, S., Streib, T., Rohrer, J. S. and Pohl, C. A. "Improved method for the determination of trace perchlorate in ground and drinking waters by ion chromatography" J. Chromatogr., A 2000, 888, 151-158. 44 Roehl, R., Slingsby, R., Avdalovic, N. and Jackson, P. E. "Applications of ion chromatography with electrospray mass spectrometric detection to the determination of environmental contaminants in water" J. Chromatogr., A 2002, 956, 245-254. 45 Salov, V. V. Y., J.; Shibata, Y.; Morita, M. "Determination of inorganic Halogen Species by Liquid Chromatography with Inductively Coupled Argon Plasma Mass Spectrometry" Anal. Chem. 1992, 64, 2425-2428. 46 Urbansky, E. T., Magnuson, M. L., Freeman, D. and Jelks, C. "Quantitation of perchlorate ion by electrospray ionization mass spectrometry (ESI-MS) using stable association complexes wih organic cations and bases to enhance selectivity" J. Anal. At. Spectrom. 1999, 14, 1861-1866. 47 Wendelken S. C., V. L. E., Coleman D. E., Munch D. J. U.S. Environmental Protection Agency Method 331.0 Revision 1.0 2005.

8

MS/MS48 were developed to detect perchlorate ion in water, but neither were developed

to identify and quantify other oxyhalides at the same time.

Measurement of perchlorate ion and other contaminants, such as chlorate and

bromate ions, in concentrated sodium hypochlorite solutions present several challenges to

the analytical chemist. The sodium hypochlorite solutions analyzed ranged from 0.35–

13%, as active chlorine, with specific conductance ranging from 51.8 mS/cm to 498

mS/cm (ionic strength ranging from 0.8-8M). The concentration differences between

hypochlorite ion and other ions of interest can differ by several orders of magnitude, thus

requiring multiple dilutions and/or multiple methods for the determination of each

analyte. Chloride, sulfate, and phosphate ions have been shown to cause potential

interference when determining perchlorate ion concentration by LC-MS/MS in water49,

which may necessitate higher dilutions or sample clean-up steps. Thus, in order to

overcome potential matrix effects and possible ionization suppression, low method

detection limits (MDL) are necessary in order to account for the high sample dilutions.

Given that other contaminants such as chlorate and bromate ions have been detected

together with perchlorate ion in sodium hypochlorite solution50, quantitation of such

analytes must also be validated.

Recently, it has been demonstrated that LC-MS/MS can be used to accurately

quantitatively measure perchlorate ion together with bromate, chlorate, and iodate ions in

a variety of sample matrices, such as bottled water51 and food supplements52. However,

these methods are generally time-consuming and have not been validated for analysis of

concentrated sodium hypochlorite solutions. Electrochemical techniques have been

validated for measuring chlorate, chlorite, and hypochlorite ions, in water and sodium

hypochlorite solutions by amperometric titration.53 Direct potentiometric titrations54 have

48 Hedrick, E. B., T.;Slingsby, R.;Munch, D. U.S. Environmental Agency Method 332.0 Revison 1.0 2005. 49 Li, Y. E., J. George "Analysis of Perchlorate in Water by Reversed-Phase LC/ESI-MS/MS Using an Internal Standard Technique" Anal. Chem. 2005, 77, 4453-4458. 50 Asami, Kosaka and Kunikane "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply". 51 Snyder, S. A., Vanderford, B. J. and Rexing, D. J. "Trace analysis of bromate, chlorate, iodate, and perchlorate in natural and bottled waters" Environ. Sci. Technol. 2005, 39, 4586-4593. 52 Snyder, S. A., Pleus, R. C., Vanderford, B. J. and Holady, J. C. "Perchlorate and chlorate in dietary supplements and flavor enhancing ingrediants" Anal. Chim. Act. 2006, 567, 23-32. 53 Clesceri, L. S. G., A. E.; Eaton, A. D. Eds. Standard Methods for the Examination of Water and Wastewater 20th ed.; American Public Health Association: Washington, DC, 1998.

9

been used, but are too time-consuming55 or are optimized for determination of a single

oxyhalide species56. Based on the information in these papers, the iodometric titration

method was chosen as a reference method for measuring the concentration of chlorate ion

in concentrated sodium hypochlorite solutions. Direct potentiometric titration with

sulfite ion was chosen for highly accurate and selective quantitation of hypochlorite ion.

2.1.2 Transition metal ions: Co2+, Cu2+, Fe3+, Mn2+, and Ni2+

Metal ion analysis and quantification was performed using US EPA method57

200.8 on an Agilent (Palo Alto, CA) 7500c Inductively Coupled Plasma-Mass

Spectrometer (ICP-MS) with an Octopole Reaction System that uses hydrogen-helium

reaction gas to remove Ar-based isobaric and polyatomic oxide interferences. Internal

standards were used to correct for matrix interferences. Dilutions ranging from 1:10 to

1:500 were used to reduce the impact of the high total dissolved solids (TDS) background

of hypochlorite ion samples. Actual method reporting limits (MRLs), based on required

dilution for each metal ion, are summarized in Table 1.

Table 1. ICP-MS MRLs (µg/L) in water and hypochlorite ion solutions

Element Reagent Water Hypochlorite Ion

Solutions*

Manganese (Mn) 1.0 25

Iron (Fe) 5.0 125

Cobalt (Co) 2.0 25

Copper (Cu) 5.0 25

Nickel (Ni) 1.0 25

*Lowest MRLs based on dilution factor required for each metal ion

54 Adam, L. C. and Gordon, G. "Direct and Sequential Potentiometric Determination of Hypochlorite, Chlorite, and Chlorate Ions When Hypochlorite Ion is Present in Large Excess" Anal. Chem. 1995, 67, 535-540. 55 Ikeda, Y., Tang, T. and Gordon, G. "Iodometric method for determination of trace chlorate ion" Anal. Chem. 1984, 56, 71-73. 56 Miller, K. G., Pacey, G. E. and Gordon, G. "Automated iodometric method for determination of trace chlorate ion using flow injection analysis" Anal. Chem. 1985, 57, 734-737. 57 Creed, J. T. B., C.A.; Martin, T.D. U.S. Environmental Agency Method 200.8 Revison 5.4 1994.

10

2.1.3 Specific Conductance, Ionic Strength, and pH Measurements Specific conductance measurements were performed using a HACH Ion-Series

Conductivity / Total Dissolved Solids meter (Hach Company, Loveland, CO). A

calibration was performed using a 1,000 µmho/cm standard prior to sample analysis. In

most cases dilutions were required to bring the conductivity of hypochlorite ion solutions

to within the meter’s linear range of 20-200,000 µmho/cm. Specific conductance was

measured in order to determine the ionic strength of hypochlorite ion solutions. Ionic

strength (I) was calculated from specific conductance (σ) based on relationship

established by the Russell approximation given by Equation 2. For convenience μ is used

as the symbol for ionic strength in the later chapters, expressed in units of mol/L.

I (mol/L) = 1.6 x 10-5 x σ (µmho/cm) (2)

Total dissolved solids (TDS) can be approximated by the Langelier approximation given

by Equation 3:

TDS (mg/L) = I (mol/L) x 4 x 104 (3)

An AP62 pH/mV Meter (Fisher Scientific, Pittsburg, PA) was used to measure

sample pH. Calibration of the pH meter was performed prior to measurements of

samples using standard pH buffers (pH 4, 7, 10). Although, measurement of pH above

11.5 as concentration of hydroxide ion, determined by titration, is more accurate, pH

values are reported for convenience.

2.2 Results and discussion

2.2.1 The LC-MS/MS Analysis of Perchlorate, Bromate, and Chlorate Ions

Tandem mass spectrometry was performed using an API4000 triple-quadrupole

mass spectrometer (Applied Biosystems, Foster City, CA) equipped with an electrospray

ionization source operated in negative ion mode. When using tandem mass spectrometry,

the ionized target analyte (precursor ion) is first selected by the first quadrupole, then the

precursor ion is fragmented in the second quadrupole (by collision-induced dissociation),

while a representative specific fragment (product ion) of the precursor ion is selected by

the third quadrupole. The monitoring of the transition(s) from precursor(s) to product

ion(s), a technique known as multiple reaction monitoring (MRM), makes analysis by

11

MS/MS very selective and sensitive58. The following precursor/product ion transitions

were used: 35ClO4- (m/z 99) to 35ClO3

- (m/z 83) for perchlorate ion; 35ClO3- (m/z 83) to

35ClO2- (m/z 67) for chlorate ion; and 79BrO3

- (m/z 127) to 79BrO2- (m/z 111) for bromate

ion. The following confirmation transitions were used: 37ClO4- (m/z 101) to 37ClO3

-(m/z

85) for perchlorate ion; 37ClO3- (m/z 85) to 37ClO2

- (m/z 69) for chlorate ion; and 81BrO3-

(m/z 129) to 81BrO2- (m/z 113) for bromate ion. Perchlorate and bromate ions were

quantified using isotope dilution. Stable isotope-labeled versions of perchlorate

(35Cl18O4) and bromate (79Br18O4) ions were used. As no source of oxygen-18 labeled

chlorate ion was commercially available, chlorate ion was quantified using external

calibration.

Perchlorate and chlorate ion standards (>99.5 % purity) were obtained from Ultra

Scientific (North Kingstown, RI) and J.T. Baker (Phillipsburg, NJ), respectively.

Bromate ion standard was obtained from Ultra Scientific (North Kingstown, RI). The

stable-isotope labeled perchlorate (Cl18O4-) and bromate (Br18O3

-) ions used as internal

standards were obtained from Icon Isotopes (Summit, NJ). Trace analysis grade

methanol was obtained from Burdick and Jackson (Muskegon, MI). Formic acid (1 mL

ampoules, 99+ %) was purchased from Thermo Scientific (Rockford, IL). Standard and

sample dilutions were done using deionized reagent water purified by Milli-Q Gradient

System (Millipore, Billerica, MA).

In order to reduce the run times, analytes were separated using a 75 x 4.6 mm

Synergi Max-RP C12 column with a 4 μm pore size (Phenomenex, Torrance, CA). An

injection volume of 20 μl was used for all samples. A binary gradient consisting of 0.1

% formic acid (v/v) in water (A) and 100 % methanol (B) at a flow rate of 700 μl/min

was used. The gradient was as follows: 2 % B held for one minute, increased linearly to

15 % B by two minutes, changed to 95 % B and held for four minutes, and finally

changed to 2 % B and held for three minutes. A one minute equilibration step at 2 % B

was used at the beginning of each run to bring the total run time per sample to ten

minutes. With the use of a shorter column, the analysis time was reduced to half of the

58 Zakett, D., Flynn, R. G. A. and Cooks, R. G. "Chlorine isotope effects in mass spectrometry by multiple reaction monitoring" J. Phys. Chem. 1978, 82, 2359-2362.

12

original method59 with all three analytes eluting after two minutes. The effects of faster

separation on method performance were further investigated.

2.2.2 Validation of LC-MS/MS Method for the Analysis of Hypochlorite

Solutions

Using the optimized separation method, standards containing 0.02, 0.1, and 0.5

μg/L standards of perchlorate, bromate, and chlorate ions were injected and analyzed 8

times. To calculate a method detection limit (MDL), standard deviation of the 8 replicate

data points was multiplied by the Student’s T value for 7 degrees of freedom. These

results are shown in Table 2.

When a 0.02 μg/L perchlorate ion standard was injected, an S/N ratio of 6.3 was

obtained, as shown in Figure 2. The relative standard deviation (RSD = Standard

Deviation / Mean x 100%) of 8 replicate measurements of the 0.02 μg/L perchlorate ion

standard is almost twice as large as that of the 0.1 μg/L perchlorate ion standard but is

still below 10% RSD. Bromate and chlorate ions MDLs of 0.06 μg/L and 0.23 μg/L were

obtained, respectively. Practical quantitation limits (PQLs) for each analyte at the

instrument were calculated by multiplying the MDL by 3 for each analyte and rounding

up (the conservative MDL of 0.015 μg/L was used for ClO4-). This resulted in PQLs of

0.05 μg/L for ClO4-, 0.20 μg/L for BrO3

-, and 0.70 μg/L for ClO3

-. Analysis of most

sodium hypochlorite solutions for measurement of perchlorate ion rarely required a

PQL< 0.1 μg/L, thus the lowest calibration standard for typical analysis contained 0.10

μg/L of ClO4-, 0.20 μg/L of BrO3

-, and 1.0 μg/L of ClO3-. To demonstrate the

applicability of the PQL to a complex sample matrix, chromatograms of sample and

standard perchlorate ion solutions are shown in Figure 3.

59 Snyder, Vanderford and Rexing "Trace analysis of bromate, chlorate, iodate, and perchlorate in natural and bottled waters".

13

Table 2. MDL data for perchlorate, bromate, and chlorate ions (n = 8) Standard Concentration

Replicate 0.02 μg/L 0.10 μg /L 0.10 μg /L 0.50 μg /L 1 0.022 0.098 0.108 0.545 2 0.023 0.106 0.086 0.422 3 0.025 0.100 0.133 0.415 4 0.023 0.101 0.121 0.331 5 0.025 0.107 0.076 0.409 6 0.027 0.099 0.122 0.498 7 0.023 0.096 0.099 0.525 8 0.027 0.110 0.112 0.361

Mean 0.024 0.102 0.107 0.438 Std. Dev. 0.002 0.005 0.019 0.077 RSD (%) 8.3 4.9 17.8 17.6

Student' t @ 98% n-1 2.998 2.998 2.998 2.998 MDL = SD X Student's t 0.006 0.015 0.058 0.231 PQL= 3 x MDL(μg /L)* 0.020 0.050 0.200 0.700 *Values rounded up

0.5 1.0 1.5 2.0 2.5 3.5 4.0 4.5Time, min

050100150200250300350400450500550600650700

Inte

nsity

, cps

2.03

S/N = 6.3

Peak Int.(Subt.)=7.0e+2

3xStd.Dev.(Noise)=1.1e+2

3.00.0

0.02 μg/L ClO4- Standard Solution

Figure 2. Extracted ion chromatogram of perchlorate ion MRM m/z 98.9/82.8 of standard solution containing 0.02 μg/L of perchlorate ion

14

200400600800

12001400

1000

2200

Inte

nsity

, cps

2.010.10 μg/L ClO4

- Sodium Hypochlorite Solution

S/N = 13.5

Peak Int.(Subt.)=2.2e+3

3xStd.Dev.(Noise)=1.6e+2

0.5 1.0 1.5 2.0 2.5 3.5 4.0 4.5Time, min

3.00.00

160018002000

0200400600800

1000

18002000220024002600

2.03

S/N = 12.9

Peak Int.(Subt.)=2.6e+33xStd.Dev.(Noise)=2.0e+2

0.10 μg/L ClO4- Standard Solution

14001600

1200

0.5 1.0 1.5 2.0 2.5 3.5 4.0 4.5Time, min

3.00.0

Inte

nsity

, cps


of (a) sodium hypochlorite solution with 0.1 μg/L ClO4-; (b) standard

solution with 0.1 μg/L ClO4-

The signal-to-noise (S/N) ratio was calculated using Analyst 1.5 Software

(Applied Biosystems) without curve smoothing. As evident from Figure 3, very

comparable S/N ratios and raw counts were observed in standard and sodium

hypochlorite sample solutions containing 0.1 μg/L ClO4-, indicating the applicability of

the method to separate matrix interferences while selectively monitoring for perchlorate

ion. Perchlorate and bromate ions were calibrated using isotope dilution with 1/x2

(a)

(b)

15

weighting and typical R2 ≥ 0.99. For chlorate ion, an external calibration with 1/x2

weighting was used with typical R2 ≥ 0.99.

In order to examine any possible impacts of matrix interferences (e.g., ionization

suppression or enhancement, isobaric interferences, and chromatographic resolution) at

different dilutions, a series of dilution tests were completed for the measurement of

perchlorate, bromate, and chlorate ions by LC-MS/MS. Due to known isobaric

interferences by bisulfate ion (HSO4-) on perchlorate ion, the use of 2.5cc OnGuard II Ba

and 2.5cc OnGuard II H Cartridge (Dionex, Sunnyvale, CA), to reduce bisulfate and

carbonate ion concentrations, was investigated.

The cartridges were conditioned by flushing and discarding 30 mL deionized

water. Sample solutions were eluted typically at a flow rate of 2.0 mL/min using a

mechanical syringe pump (KD Scientific, Holliston, MA) and at least the first 10 mL of

eluent were discarded prior to collecting the sample aliquot for analysis. Split aliquots of

the dilution test samples were sequentially passed through each type of cartridge. The

results of dilutions on measured analyte concentration with and without the clean-up step

are summarized in Table 3.

There were no major differences observed for any of the oxyhalide analytes

between filtered and unfiltered samples. Thus, it was decided that if bisulfate ion

interference was observed, the use of a clean-up/filtration step with OnGuard II Ba

cartridges could be employed without negatively impacting the analysis.

Table 3. Spike recoveries of analytes with and without filtration and at different dilution factors (DF) (n = 3)

ClO4- BrO3

- ClO3-

Dilution Factor (DF) 5000 1000 100 10 5000 1000 100 10 5000

% Recovery 99 99 96 95 100 94 86 7.7* 91 Std. Dev. 3.2 2.3 4.6 5.1 2.3 2.5 17.2 0.8 11.4 N

o Fi

ltrat

ion

RSD (%) 3.3 2.4 4.8 5.4 2.2 2.7 20 10.7 12.5 % Recovery 107 100 100 100 99 90 89 -16* 98

Std. Dev. 2.3 3.5 2.3 3.9 5.9 6.8 3.5 40.8 10.2 Ba/

H

Filtr

atio

n

RSD (%) 2.1 3.3 2.3 3.9 6 7.5 4 250 10.5 MRL =PQL X DF (μg/L) 250 50 5 0.5 1,000 200 20 2 3,500 *Dilution is inadequate resulting in concentrations impeding analysis

16

The analysis of perchlorate ion showed recoveries of 99.5 % (± 3.6) for a wide

range of dilutions, analyte concentrations, and Ba/H clean-up steps. Perchlorate ion could

be quantified at dilutions as low as 1:10, thereby achieving MRL as low as 0.5 μg/L

(dilution factor times the PQL) using the current method. Analysis of bromate ion was

the most susceptible to the matrix effects. Examples of bromate ion chromatograms at

different dilution factors are shown in Figure 4.

Figure 4. Bromate ion chromatograms of sodium hypochlorite sample diluted

by: (a) factor of 1:10; (b) factor of 1:100; (c) factor of 1:1000

79Br18O3 79Br16O3

79Br18O3 79Br16O3

(a)

(b)

(c)

79Br18O3 79Br16O3

17

Interestingly, the loss of 18O-labeled bromate ion signal was much higher than

that of the analyte signal at 1:10 and 1:100 dilutions, thereby producing erroneous results

(figure 3). A higher dilution of 1:1000 eliminated this problem since the peak-widening

was significantly reduced and showed no effect on the accuracy based on recovery data

shown in Table 3. In most cases, matrix interferences were minimized by diluting the

samples, while the sensitivity of the LC-MS/MS method still provided low MRLs (Table

3). Thus, the observed signal suppression was resolved simply by analyzing samples at

higher dilutions. Figure 5 illustrates the improvement in accuracy based on the dilution

factor.

0

2

4

6

8

10

10 100 1000 5000Dilution Factor

mg/

L C

lO4-

and

BrO

3-

0

50

100

150

200

250

300

350

mg/

L C

lO3-

Bromate

PerchlorateChlorate

Figure 5. Actual sample concentrations of analytes measured at different

dilutions

In “freshly prepared” hypochlorite ion solutions where chlorate ion concentration

was still low, a single dilution could be used to measure all three oxyhalides. However,

as sodium hypochlorite solutions age, chlorate ion forms, thus higher dilutions may be

needed for the analysis of older solutions, thereby requiring more than one dilution per

sample. Carrying out serial dilutions of several orders of magnitude may compound

errors associated with each dilution resulting in higher variability. Alternatively,

determinations of higher concentrations of chlorate ion can also be performed by

iodometric titration, thereby eliminating the need to analyze multiple dilutions using LC-

MS/MS.

18

2.2.3 Iodometric Titrations: Analysis of Chlorite, Chlorate, and

Hypochlorite Ions

Many variations of iodometric titration methods have been developed over the

years and rely on the principle that hypochlorite, chlorite, and chlorate ions react with

iodide ion (I-) to produce triiodide ion (I3-), in presence of excess iodide ion, shown by

Equations 4-6:

OCl- + 3I- + 2H+ → I3- + Cl- + H2O pH 1.3 (4)

ClO2- + 6I- + 4H+→ 2I3

- + Cl- + 2H2O pH 1.3 (5)

ClO3- + 9I- + 6H+→ 3I3

- + Cl- + 6H2O 6M H+ (6)

Triiodide ion can be titrated with thiosulfate ion (S2O32-) or sulfite ion (SO3

2-),

shown by Equations 7 and 8, thus allowing determination of these oxyhalide anions.

I3- + 2S2O3

2- → 3I- + S4O62- (7)

I3- + SO3

2- + 2OH- → SO42- + 3I- + H2O (8)

However, since both chlorite and hypochlorite ions react with iodide ion at pH

1.3, measuring concentrations of chlorite and hypochlorite ions separately requires an

external technique to measure one of the analytes. However, the reaction of chlorate ion

with iodide ion takes place under highly acidic conditions ([H+] = 6 M), requiring the use

of concentrated hydrochloric acid (HCl). Thus, measurement of chlorate ion is easily

separated from measurement of chlorite and hypochlorite ions.

A well established way to make a standard solution of triiodide ion is to add a

known amount of Reagent Grade iodate ion (IO3-) to an acidic solution containing a small

excess of iodide ion, shown in Equation 10. Thus, the prepared standard solution of

triidodie ion can be used to standardize thiosulfate ion and sulfite ion solutions.

IO3- + 8I- + 6H+ ↔ 3I3

- + 3H2O (9)

Similarly, bromate ion reacts with iodide ion, shown by Equation 10.

BrO3- + 9 I- + 6 H+ →3 I3

- + 3 H2O pH 1.3 (10)

However, if other ions such as hypochlorite and chlorite ions are present,

differentiation of analytes is not possible. Thus, for experiments involving bromate ion

and chlorite ion spikes, measurement of hypochlorite ion by iodometric titration would

require a separate analysis to determine concentrations of bromate and chlorite ions.

19

2.2.3.1 Adam-Gordon Method

Direct potentiometric titration of hypochlorite ion with sulfite ion allows selective

determination of hypochlorite ion from chlorite ion and bromate ions.60 This

potentiometric titration technique, combined with iodometric determination of chlorate

ion allows selective determinations of specific anions. Thus, hypochlorite ion solutions

can be analyzed for all three analytes in sequential steps, outlined below and based on

equations 11, 9, 8, 6, and 5:

Step 1: Determination of hypochlorite ion:

(a) Adjust pH to 10.5, using 0.4 M borate buffer

(b) Titrate hypochlorite ion with sulfite ion, given by equation 11:

OCl- + SO32- → SO4

2- + Cl- pH 10.5 (11)

(c) To remove any excess sulfite ion, add potassium iodide, then

triiodide ion solution. Triiodide ion reacts with sulfite ion,

given by equation 8:

I3- + SO3

2- + 2OH- → SO42- + 3I- + H2O (9)

(d) Remove any excess triiodide ion by adding dilute thiosulfate

ion solution:

I3- + 2S2O3

2- → 3I- + S4O62- (8)

Step 2: Determination of chlorite ion:

a) Adjust pH to 1.3 using 3 M HCl, reaction 4 takes place:

ClO2- + 6I- + 4H+ → 2I3

- + Cl- + 2H2O (5)

b) Back-titrate triiodie ion with thiosulfite ion:

I3- + 2S2O3

2- → 3I- + S4O62- (8)

Step 3: Determination of chlorate ion:

a) Purge sample solution with nitrogen gas at least 5 minutes

b) Add concentrated HCl, reaction 5 takes place:

ClO3- + 9I- + 6H+ → 3I3

- + Cl- + 3H2O 6M H+ (6)

b) Back-titrate triiodie ion with thiosulfite ion:

I3- + 2S2O3

2- → 3I- + S4O62- (8)

60 Adam and Gordon "Direct and Sequential Potentiometric Determination of Hypochlorite, Chlorite, and Chlorate Ions When Hypochlorite Ion is Present in Large Excess".

20

Titrations with sulfite ion were performed using a VIT 90 Video Titrator with a

P101 platinum k401 SCE electrode pair (Radiometer, Copenhagen, Denmark). Titrations

with thiosulfate ion were performed with a standard 50 mL laboratory glass burette.

Borate buffer (0.4 mol/L), was prepared using boric acid (ACS Reagent, 99.5 %), bought

from Alfa Aeser (Ward Hill, MA); and sodium hydroxide (ACS Reagent, 97%), obtained

from VWR (West Chester, PA). Borate buffer was used to adjust hypochlorite ion sample

solutions to pH 10.5. Hydrochloric acid (ACS Reagent, 37 %), obtained from Sigma-

Aldrich (St. Louis, MO) or Fisher Scientific (Pittsburg, PA), and the sample solution

were purged with nitrogen gas to minimize oxidation of iodide ion by oxygen prior to

chlorate ion determination.

To adjust hypochlorite ion sample solutions to pH 1.3, a 3 mol/L HCl solution

was used. Potassium iodate (99.4-100.4 % ACS Reagent) and potassium iodide (99 %)

were obtained from Sigma Aldrich and Alfa Aeser (Ward Hill, MA). Sodium sulfite (98

% ACS Reagent) and sodium thiosulfate (99 %, ACS Reagent) were obtained from

Sigma-Aldrich (St. Louis, MO). Dilute (≤ 10 mmol/L) thiosulfate ion solution was

prepared and used for removal of excess triiodie ion during the determination of

hypochlorite, chlorite and chlorate ions. Potassium iodate standard (0.1 M) was prepared

weekly and used to standardize sulfite and thiosulfate ion solutions daily. A solution

containing ~0.02 M I3- was prepared by adding an aliquot of ~0.1 M IO3

- standard to

acidic solution (pH ~1) containing excess iodide ion (~0.5 M I-) and used to remove

sulfite ion. Standardization of 0.2 M SO32- and 0.1 M S2O3

2- solutions by 0.1 M IO3-

solution resulted in standard deviations of less than three parts per thousand and less than

1% relative standard deviation. Table 4 shows results of typical standardization of sulfite

and thiosulfate ion solutions.

Table 4. Standardization of sulfite and thiosulfate ion solutions by 0.109 M IO3-

mL of IO3-

Standard mL SO32- [SO3

2-] mL of IO3

- Standard mL S2O3

2- [S2O32-]

0.500 0.6837 0.2399 1.000 8.50 0.0772 0.700 0.9505 0.2416 1.000 8.40 0.0781 0.700 0.9539 0.2408 1.000 8.50 0.0772

Mean 0.241 Mean 0.0775 Std. Dev. 0.0009 Std. Dev. 0.0005 RSD (%) 0.353 RSD (%) 0.685

21

2.2.4 Method Selection for the Measurement of Chlorate Ion

Sodium hypochlorite solutions are obtained as either bulk commercial bleach as

3-13 % free available chlorine (FAC), or as on-site generated (OSG) sodium hypochlorite

solutions that can range 0.3 – 3 % as FAC. Bulk sodium hypochlorite solutions, 10–13%

FAC, were obtained from Acros Organics USA (Morris Plains, NJ) and VWR (Brisbane,

CA). OSG sodium hypochlorite solution, ~0.8 % FAC, was obtained from a sodium

hypochlorite on-site generator at River Mountains Water Treatment Facility (RMTF)

(Henderson, NV). To determine which method (iodometric titration or LC-MS/MS) was

most reproducible for chlorate ion analysis of either bulk or OSG sodium hypochlorite

solutions, replicate samples were split and analyzed by both methods and the results

compared. Pure sodium chlorate (ACS Grade, ≥99 %) was purchased from VWR

(Brisbane, CA) and used to spike sodium hypochlorite solutions and prepare chlorate ion

standard solutions. A comparison of the two methods for analysis of seven replicate bulk

hypochlorite ion solutions is shown in Table 5.

Table 5. Comparison of measurements by the LC-MS/MS and iodometric titration for bulk hypochlorite ion solutions (n = 7)

Non-spiked Hypochlorite ClO3- Spiked % Recovery

Replicate LC-MS/MS* Titration* LC-MS/MS* Titration* LC-MS/MS* Titration* 1 16.9 17.7 24.4 25.2 95.7 95.1 2 18.0 17.9 24.4 24.8 81.7 88.3 3 15.9 17.7 23.5 25.2 97.0 95.1 4 15.4 17.7 23.1 25.0 98.3 92.8 5 16.1 17.9 25.1 25.2 115 92.8 6 12.9 17.9 24.0 25.2 142 92.8 7 16.5 18.1 23.0 25.4 83.0 92.8

Mean 16.0 17.9 23.9 25.1 102 92.8 Std. Dev. 1.6 0.1 0.8 0.2 20.8 2.3 RSD (%) 9.9 0.8 3.2 0.7 20.4 2.4 *All concentration data listed as g/L

These results indicate that, although both methods produce similar results,

titration of the concentrated hypochlorite ion solutions resulted in much lower relative

standard deviations (RSD) than the LC-MS/MS. This higher variability observed in the

LC-MS/MS data is most likely a result of carrying out serial dilutions of several orders of

22

magnitude, which may have compounded the errors associated with each dilution. For

example, in order to analyze hypochlorite ion solutions that contain chlorate ion at 100

g/L, a dilution factor of at least 400,000 would be needed to bring the concentration of

chlorate ion to the middle of a typical calibration curve of 1 – 500 μg/L. The effect of

this variability on analysis of samples during a chlorate ion spike experiment can be

visually observed from the comparison of the results from two methods of duplicate

samples, shown in Figure 6 (for illustration purposes, smooth lines were used to follow

changes in chlorate ion concentration as a function of time).

In several chlorate ion spiked experiments, dilutions on the order of 1:2,500,000

were required. The error bars of chlorate ion measurements by LC-MS/MS method,

based on analysis of duplicate samples solutions, were much higher for samples spiked

with higher chlorate ion concentrations, than those obtained by iodometric titration, as

shown in Figure 6.

However, given that the standard deviation (SD) of ± 0.2 g/L (Table 5) or ± 0.002

M, the iodometric titration method should be used to measure concentration of chlorate

ion above 2.0 g/L (SD=0.2 g/L X 10) or 0.025 M to achieve reasonable precision (RSD <

10 %). The results shown in Table 6 indicate poor precision when measuring

concentration of chlorate ion less than 1.0 g/L by iodometric titration.

Table 6. Comparison of measurements by the LC-MS/MS and iodometric titration for OSG sodium hypochlorite solutions at less than 1.0 g/L ClO3

- (< 10 mM) (n ≥ 3) Titration LC-MS/MS

Mean (g/L ClO3-) 0.56 0.31

Std. Dev. (g/L ClO3-) 0.20 0.01

RSD (%) 35.2 1.55 (n=4) (n=3)

23

0

50

100

150

200

250

0 2 4 6 8 10Days

g/L

ClO

3-

Control Control + 50g/L ClO3

-

Control + 100g/L ClO3-


0

50

100

150

200

250

0 2 4 6 8 10Days

g/L

ClO

3-

Control Control + 50g/L ClO3

-



Figure 6. Comparison of chlorate ion measurements by (a) iodometric titration and

(b) by the LC-MS/MS, during a chlorate ion spike experiment, at 75 ºC (Control = solution at initial [OCl-] = 1.46 M and [ClO3

-] = 0.29 M)

Although accurate and reproducible trace analysis of chlorate ion by iodometric

titration is possible61, it was simply not necessary. All of the chlorate ion spiked

experiments, which were used for kinetic determinations, were performed with the

iodometric titration method described herein, which provided the necessary precision, as

evident from results shown in Table 5 and Figure 6. To ensure reliability of the data and

to minimize variation associated with the analysis, the iodometric titration method was

used for the determination of chlorate ion concentrations in 10 – 250 g/L range (generally

61 Ikeda, Tang and Gordon "Iodometric method for determination of trace chlorate ion".

(a)

(b)

24

in hypochlorite ion solutions with > 5 % FAC i.e. bulk sodium hypochlorite), which

included chlorate ion spike experiments. When analyzing samples that required dilution

factor of less than 1:100,000, chlorate ion measurements by LC-MS/MS method,

typically had an RSD of less than 5 % (n = 3). Thus the LC-MS/MS method can be used

reliably for measurement of chlorate ion in the 0.01 – 10,000 mg/L range, but was not

required for chlorate ion spike experiments.

2.2.5 Selection of Quenching Agent

Refrigerating sample aliquots at 4 ºC significantly slows the decomposition of

hypochlorite ion: the half life of 13% sodium hypochlorite solution at 25 ºC is 130 days;

at 4 ºC the half life is 3184 days, according to the Bleach 2001. However, for any

experiments involving the measurements of concentration of perchlorate ion in

hypochlorite ion sample solutions collected in the field, it was not always possible to

attenuate the decomposition reaction rate of hypochlorite ion. In order to stop the

formation of perchlorate ion in sodium hypochlorite solutions over time and thus preserve

the integrity of the sample for analysis, a quenching agent was needed to neutralize

hypochlorite ion. Thus, quenching agents were considered for sample preservation and

also to reduce instrument maintenance.

A total of seven hypochlorite ion quenching agents were investigated. They were

ascorbic acid, malonic acid, oxalic acid, glycine, sodium sulfite, sodium thiosulfate, and

hydrogen peroxide. A brief summary of the known mechanisms and operating ranges for

each of the quenching agents is summarized below and is based on the Doctoral

Dissertation Research performed by Wood62, Sweetin63, and Adam64 at Miami

University. For example, ascorbic acid was found to be most effective over the pH range

of 3-11, glycine above pH 8.5, oxalic acid below pH 8.5, malonic acid over the pH range

of 3-8.5. A selective reaction of sulfite ion and hypochlorite ion at pH 10.5 led to the

62 Wood, D. W. I. Determination of Disinfectant Residuals in Chlorine Dioxide Treated Water Using Flow Injection Analysis Ph.D. Dissertation, Miami University, Oxford, OH, 1990. 63 Sweetin, D. L. Developments in the Analytical Methodology For the Determination of Free available Chlorine, Inorganic Chloramines, and Oxyhalogen Species Ph.D. Dissertation, Miami University, Oxford, OH, 1993. 64 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

25

development of an analytical method for the accurate determination of hypochlorite ion65,

Equation 11:

OCl- + SO32- → SO4

2-+ Cl- pH 10.5 (11)

At pH 10.5 sulfite ion will only react with hypochlorite and not chlorite or

chlorate ions. Similarly, thiosulfate ion selectively reacts with hypochlorous acid at pH

6.3, Equation 12:

S2O32- + 4HOCl + H2O → 2SO4

2-+ 4Cl-+ 6H+ pH 6.3 (12)

Hydrogen peroxide can act as either a reducing agent or oxidizing agent,

depending on the pH. In the case of reaction with hypochlorite ion in basic solution it acts

as a reducing agent, Equation 13.

OCl- + H2O2 → Cl- + H2O + O2 (13)

Hydrogen peroxide is also known to react with iodide ion66 as well as sodium

thiosulfate and sodium sulfite,67 as given by Equations 14-16. Thus, any excess hydrogen

peroxide from the reaction with hypochlorite ion can potentially interfere with the

measurement of chlorite and chlorate ions.

3I- + H2O2 + 2H+ → I3- + 2H2O (14)

2S2O32- + H2O2 → S4O6

2- + 2OH- (15)

SO32- + H2O2 → SO4

2- + H2O (16)

To guide the selection of a hypochlorite ion quenching agent, several deciding

criteria were identified:

• Safety, ease of handling, transport, and stability

• Ability to quench hypochlorite ion reproducibly

• Impact on perchlorate ion levels (LC-MS/MS method)

• Impact on bromate ion levels (LC-MS/MS method)

• Impact on chlorate ion levels (LC-MS/MS and iodometric titration)

65 Adam and Gordon "Direct and Sequential Potentiometric Determination of Hypochlorite, Chlorite, and Chlorate Ions When Hypochlorite Ion is Present in Large Excess". 66 Liebhafsky, H. A. and Mohammad, A. "A Third-order Ionic Reaction without Appreciable Salt Effect" J. Phys. Chem. 1934, 38, 857-866. 67 Liu, W., Andrews, S. A., Stefan, M. I., & Bolton, J. R. "Optimal methods for quenching H2O2 residuals prior to UFC testing" Water Res. 2003, 37, 3697-3703.

26

Ascorbic acid, USP Grade, was obtained from Mallinckrodt Chemicals

(Philipsburg, NJ). Glycine, ACS Reagent 98.5%, was supplied by Sigma Aldrich.

Hydrogen peroxide, 30-32%, Assay 99.999% min., was obtained from GFS Chemicals

(Columbus, OH). Malonic acid, 99 %, was purchased from Acros Organics USA (Morris

Plains, NJ). Oxalic acid dehydrate, ACS Grade, was obtained from EMD Chemicals

(Gibbstown, NJ). Sodium sulfite and sodium thiosulfate used for the titration methods

described in this chapter were also used in the quenching experiments.

2.2.5.1 Safety, Ease of Handling, Transport, and Stability

An initial ratio of 1 mole of quenching agent to 1 mole of hypochlorite ion was

used for each test. The quenching agent was pre-weighed (in case of oxalic acid, an

aliquot of stock solution was used) and placed in a high density polyethylene (HDPE)

sample container in a chemical hood. An aliquot of 10 mL of 13 % NaOCl solution was

added. The solution was stirred to dissolve the quenching agent and allowed to react for

at least 5 minutes. An aliquot of the sample was tested for FAC residual using Hach

DPD test kit. Several of the quenching agents under consideration (hydrogen peroxide,

glycine, and oxalic acid) reacted vigorously with hypochlorite ion, causing in some cases,

significant loss of sample during the reaction and/or produced heat and noxious fumes.

Concentrated solutions of hydrogen peroxide (32% w/w) required special

handling, due to the hazardous nature of H2O2. Quenching 10 mL of 13 % sodium

hypochlorite solution produced a considerable amount of heat and gas, making this

quenching agent the most dangerous to use. Glycine, though relatively safe itself, reacted

violently and produced a very noxious gas when reacted with hypochlorite ion.

Even though all of the test reactions were performed in a well-ventilated chemical

hood, the gas produced from reaction between glycine and hypochlorite ion caused

feelings of light-headedness, dizziness, and nausea. Thus, the use of glycine would

require special precautions when collecting samples in the field. The remaining

quenching agents also produced heat and gas, but to a lesser extent. Ascorbic and

malonic acids reacted the least vigorously and appear to be the safer and easier-to-handle

quenching agents tested. None of the quenching agents had associated

27

shipping/transportation restrictions (other than including MSDS information with each

shipping carton), with the exception of concentrated hydrogen peroxide.

Regarding stability, hydrogen peroxide (32% w/w) solution has limited stability

and a limited shelf-life. Ascorbic acid produced marked color changes that can interfere

with titrimetric analyses, both during storage of a 1 M stock solution (Figure 7) and after

quenching of utility hypochlorite samples. Stock solution color change ranged from a

colorless solution upon first preparation, to a yellow solution after 20 days of storage at

ambient temperature, to a dark red solution after 37 days of storage. Quenched

hypochlorite ion solutions exhibited similar color changes due to presence of excess

ascorbic acid (Figure 8). The development of color in quenched concentrated

hypochlorite ion samples would interfere with the determination of chlorate ion by the

iodometric titration.

Figure 7. Stock solutions of ascorbic acid freshly prepared (left), after 20 days

(center), and after 37 days of storage (right)

28

Figure 8. Ascorbic acid -quenched hypochlorite ion sample solutions (left 3

bottles) and malonic acid-quenched solution (right bottle)

2.2.5.2 Ability to Quench Hypochlorite Ion Reproducibly

Potentiometric titration with sulfite ion was used to determine the exact molarity

of hypochlorite ion solutions prior to quenching and to calculate the mole ratio of

quenching agent to hypochlorite ion. When using sulfite ion solution for quenching

experiments, standardization with iodate ion were performed prior to use in order to

determine the correct volume required to deliver the appropriate moles of quenching

agent. All quenching agents tested were able to quench hypochlorite ion in test samples,

though oxalic acid had to be pre-dissolved in reagent grade water prior to quenching due

to its low solubility in concentrated hypochlorite ion solutions. Ascorbic acid, malonic

acid, and oxalic acid were able to quench hypochlorite ion in different concentration and

volume solutions of hypochlorite ion.

The mole ratio of quenching agent to hypochlorite ion was 1.2 for ascorbic acid,

0.75 for malonic acid, and 1.5 for oxalic acid. Glycine was the least reproducible in

quenching hypochlorite ion with mole ratios varying from 0.20 to 0.53; one hypothesis is

that the reaction between glycine and hypochlorite ion is temperature and pH dependent

and thus may produce variable results depending on sample composition and handling.

Oxalic and malonic acids were the slowest reactants, requiring up to one hour for

complete quenching of 13 % hypochlorite ion solution. Sodium thiosulfate and sodium

sulfite were found to quench hypochlorite ion reproducibly, and sodium sulfite was used

routinely to determine the concentration of hypochlorite ion.

29

2.2.5.3 Impact on Analysis of Bromate, Chlorate, and Perchlorate Ion

Ascorbic acid, sodium thiosulfate and sodium sulfite were observed to have an

adverse effect on the analysis of bromate ion by LC-MS/MS. Both bromate ion and the

bromate ion internal standard were negatively impacted by the presence of thiosulfate and

sulfite ions. The chromatograms of sulfite-quenched and thiosulfate-quenched samples,

shown in Figure 9, illustrate the peak shifts, peak attenuation, and peak splitting (Figure 4

(c), page 16 can be used as reference).

Figure 9. Chromatogram of bromate (left) and 18O-labeled bromate (right) of

(a) sulfite-quenched sample, and (b) thiosulfate-quenched sample of 13% sodium hypochlorite solution diluted by a factor of 1:10,000

Furthermore, suppression of both analyte and internal standard signals hindered

the ability for adequate correction by isotope-dilution, resulting in poor recoveries for

samples quenched with thiosulfate ion (60%, n = 3) and no quantifiable recoveries for

(a)

(b)

79Br18O3 79Br16O3

79Br18O3 79Br16O3

30

samples quenched with sulfite. Ascorbic acid also negatively impacted analysis of

bromate ion. When hypochlorite ion solutions, quenched with ascorbic acid were

analyzed, no detectible bromate ion was observed. Furthermore, spiked ascorbic acid-

quenched hypochlorite ion solutions with bromate ion standard, showed much lower

recoveries (49%, n=3) than the non-quenched hypochlorite ion solutions, indicating that

excess ascorbic acid in fact may reduce bromate ion. Similarly, sulfite ion is also known

to reduce bromate ion in aqueous solutions.68 Malonic acid was found to have no impact

on bromate, chlorate, and perchlorate ions analysis. Recovery data obtained by LC-

MS/MS analysis of quenched and non-quenched hypochlorite ion samples are shown in

Table 7, while comparisons of quenched and non-quenched hypochlorite ion samples by

the analyte at different dilutions are shown in Tables 8.

Table 7. Effects of malonic acid (MA) on recoveries of chlorate, perchlorate, and bromate ions measured by LC-MS/MS (n = 3, replicate samples analyzed in triplicate; S.D. = standard deviation)

% Recovery (Mean ± S.D.)

Perchlorate Spike μg/L Non-quenched

sample MA quenched

sample 2.1 101.0 ± 1.8 104.0 ± 2.3 41.6 94.3 ± 1.3 94.8 ± 1.2

Bromate Spike μg/L 1.6 94.3 ± 6.5 96.7 ± 1.1

Chlorate Spike μg/L 146.6 90.8 ± 4 88.9 ± 0.5

Table 8. Effects of malonic acid (MA) on analysis of perchlorate and bromate

ions at different dilutions (n = 3; S.D. = standard deviation) Sample ClO4

- mg/L ± S.D. Sample BrO3- mg/L ± S.D.

Dilution Non-quenched

sample MA quenched

sample Non-quenched

sample MA quenched

sample 1:100,000 6.9 ± 0.6 6.7 ±0. 4 14.7 ± 2.1 13.8 ± 0.9 1:10,000 7.2 ± 0.2 6.5 ± 0.2 16.0 ± 0.3 14.8 ± 0.1

1:500 7.4 ± 0.1 7.1 ± 0.5 9.6 ± 0.6 12.9 ± 0.3

68 Keith, J. D., Pacey, G. E., Cotruvo, J. A. and Gordon, G. "Experimental results from the reaction of bromate ion with synthetic and real gastric juices" Toxicology 2006, 221, 225-228.

31

Table 9 shows measured concentrations of chlorate ion by LC-MS/MS and

iodometric titration in non-quenched sodium hypochlorite and in malonic acid (MA)

quenched sample. The difference in results and recoveries between quenched and non-

quenched samples for all three analytes is relatively small, thus malonic acid can be used

as a preservative for sodium hypochlorite ion solutions.

Table 9. Effects of quenching agent on analysis of chlorate comparison of LC-MS/MS and titration results (n = 3, replicate samples analyzed in triplicate; S.D. = standard deviation)

Sample ClO3- g/L ± S.D.

Non-quenched Hypochlorite

Malonic Acid

LC-MS/MS 16.10 ± 0.15 15.70 ± 0.15 Iodometric Titration 18.10 ± 0.23 18.00 ± 0.20 % Difference 10.8 12.5

To determine analyte stability, when using malonic acid as a quenching agent,

concentration of analytes was measured immediately after quenching and two months

later. The results showed less than 10% difference on average (based on 8 different

hypochlorite ion samples) in concentration of perchlorate, bromate, and chlorate ions

after 2-month storage, shown in Tables 10-12.

Table 10. Perchlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2-month period

μg/L ClO4-

Sample Index 4/2/2009 6/2/2009 % Difference 1 6.64 6.01 9.6 2 38.7 40.7 5.0 3 30.8 27.5 11 4 20.7 20.2 2.4 5 78.5 79.5 1.2 6 664 611 8.0 7 3100 2610 16 8 19.1 18.6 2.7 Mean 7.0 ± 5.0

32

Table 11. Bromate ion stability in hypochlorite ion solutions quenched with malonic acid over 2-month period mg/L BrO3

- Sample Index 4/2/2009 6/2/2009 % Difference

1 1.71 1.60 6.5 2 2.05 1.98 3.4 3 1.12 1.21 8.0 4 0.54 0.50 7.5 5 0.32 0.25 20 6 0.10 0.08 16 7 5.25 5.86 12 8 0.13 0.14 9.2

Mean 10.0 ± 5.3 Table 12. Chlorate ion stability in hypochlorite ion solutions quenched with

malonic acid over 2-month period mg/L ClO3

- Sample Index 4/2/2009 6/2/2009 % Difference

1 241 229 4.8 2 1025 1070 4.4 3 219 193 12 4 177 169 4.5 5 606 522 14 6 181 161 11 7 507 465 8.3 8 446 366 18

Mean 9.7 ± 5.0

In addition to the investigation of effects of various quenching agent on analysis

of bromate, chlorate, and perchlorate ions, quenched hypochlorite ion solutions were

analyzed for Co2+, Cu2+, Fe3+, Mn2+, and Ni2+ by ICP-MS method against the non-

quenched sample. No significant differences were observed, as most of metals ions were

already present at very low concentrations.

2.2.5.3 Quenching Agent Selection Summary

As a result of investigation of various effects of quenching agents on the analysis

of hypochlorite ion solution and the associated issues with the use of particular quenching

agent, Table 13 was prepared as a summary of selection criteria. Text in bold font style,

in Table 13, highlights the reasons for rejecting a given quenching agent. As such

33

glycine, hydrogen peroxide, and ascorbic acid were not recommended for quenching

based on safety, ease of handling, transport, and stability issues. Ascorbic acid and

sodium thiosulfate were not recommended for quenching due to negative impacts on

bromate ion analysis. Based on limited solubility of oxalic acid in bulk sodium

hypochlorite ion solutions, oxalic acid was not recommended for quenching. Thus,

malonic acid was chosen as the quenching agent of choice for experiments requiring

preservation (typically for samples collected off-site and requiring shipping).

Table 13. Summary of quenching agent test results and decision-making matrix Quenching Agent (QA)

Ascorbic

Acid Glycine Malonic

Acid Sodium

Thiosulfate Sodium Sulfite

Oxalic Acid

Hydrogen Peroxide

Stoichiometric Ratio (mol QA / mol OCl-)

1.2 0.55 0.75 0.28 1.1 1.5 1.1

Compatible with Iodometric Titration?

No Yes Yes Yes Yes Yes No

Effect on [BrO3-] Decrease No

Effect No

Effect Decrease Decrease No Effect

No Effect

Effect on [ClO4-] No

Effect No

Effect No

Effect No

Effect No

Effect No

Effect No

Effect

Effect on [ClO3-] No

Change No

Change No

Change No Change No Change

No Change

No Change

Solution Stable over Time No Yes Yes Yes No Yes No

Safety Issues? No Noxious

Gas, Hazard!

No No No Heat, Gas

Evolved

Violent Reaction, Hazard!

Soluble in Concentrated NaOCl ?

Yes Yes Yes Yes Yes No Yes

34

2.3 Conclusions

The optimized LC-MS/MS method was shown to meet the objective of producing

accurate and reproducible results for quantifying perchlorate, bromate, and chlorate ions

in sodium hypochlorite solutions, offering conservative PQLs of 0.05 μg/L for ClO4- and

0.20 μg/L for BrO3-. The advantage of this method is two-fold: analysis time of 10

minutes per sample and simultaneous analysis of at least two analytes (perchlorate and

bromate ions) in sodium hypochlorite solutions. This allowed increases in sample

throughput and reduced instrument maintenance.

The observed matrix interference that impacted the 18O-labeled bromate ion

internal standard was resolved by higher dilutions. Inadequate dilutions could result in

false concentrations given the differing matrix effects on the labeled and unlabeled ions.

Due to lack of availability 18O-labeled chlorate ion internal standards to correct for matrix

inferences and ionization suppression, the precision of chlorate measurements must be

monitored, especially at higher dilutions.

The iodometric titration was chosen for determinations of chlorate ion

concentrations in bulk hypochlorite ion solutions and during chlorate ion spike

experiments. The LC-MS/MS method was used for chlorate ion analysis for samples with

≤ 10 g/L ClO3-, which constituted the majority of OSG hypochlorite ion solutions. By

using both methods, the concentration of chlorate ion can be measured over a wide

concentration range, with an LC-MS/MS method PQL of 0.7 μg/L (8 nM), and

iodometric titration determinations of concentrations of chlorate ion up to 210 g/L (2.5

M) in sodium hypochlorite solutions.

Malonic acid was chosen for quenching hypochlorite ion solutions, thus

preserving the levels of perchlorate and chlorate ions based on the reproducibility to

quench, ease of handling, and minimum impacts on LC-MS/MS and titrimetric methods.

Furthermore, it was determined that using OnGuard II Ba and OnGuard II H cartridges

for sample treatment showed no major differences between filtered and unfiltered

samples. Thus if bisulfate ion is present, the use of a clean-up/filtration step with Barium

Cartridges to remove the interference with the perchlorate ion determination, does not

negatively impact analysis by the LC-MS/MS method.

35

The sodium hypochlorite samples, pending analysis, were temporary stored at 4

ºC (typically samples were analyzed within one-week period). Alternatively, malonic

acid was used during collection of hypochlorite ion solutions off-site, and water samples

that contained residual FAC.

Thus, a robust, sensitive LC-MS/MS method was developed and used for

determinations of perchlorate, bromate, and chlorate ions. Potentiometric titration with

sulfite ion was used for determination of hypochlorite ion and iodometric titration method

for determination of chlorate ion concentration in bulk hypochlorite ion solutions. Proper

storage and sample preservation conditions were determined.

36

CHAPTER 3. EXPERIMENTAL DESIGN: IDENTIFYING FACTORS

IMPACTING PERCHLORATE ION FORMATION IN

HYPOCHLORITE ION SOLUTIONS

The main objective in this section of the study was to uncover factors impacting

perchlorate ion formation in hypochlorite ion solutions. Given that perchlorate ion was

found to occur in bulk (>3.0 % FAC) sodium hypochlorite solutions and to increase over

time69 and the fact that aged sodium hypochlorite solutions containing high

concentrations of chlorate ion also had higher perchlorate ion concentrations70, it can be

inferred that perchlorate ion is forming over time and most likely a product of

hypochlorite ion decomposition. Thus, in order to address the main objective and to

design experiments, hypochlorite ion decomposition needs to be understood, and the

factors affecting the rate of decomposition considered.

Sodium hypochlorite solutions that are provided as bulk (typical commercial

bleach) are 3-13 % as FAC and have pH above 11.5. The on-site generated (OSG)

sodium hypochlorite solutions range 0.3 – 3 % as FAC and typically have a pH of 9.5.

Decomposition of hypochlorite ion and formation of chlorate ion (a product of

hypochlorite ion decomposition) in concentrated hypochlorite ion solutions, in the pH 11-

14 range, is second-order in hypochlorite ion and is at a minimum in the pH 12-13

range71. Decomposition of hypochlorite ion is acid-catalyzed and becomes more rapid at

lower pH. In the pH region of 5-8, decomposition of hypochlorite ion is a third-order

process.72 Given, that the pKa of hypochlorous acid73 at 25 ºC is 7.54, the change in the

kinetics of hypochlorite ion decomposition is attributed to the proportion of hypochlorite

ion to form hypochlorous acid.

69 Greiner, Mclellan, Bennet and Ewing "Occurrence of perchlorate in sodium hypochlorite". 70 Asami, Kosaka and Kunikane "Bromate, chlorate, chlorite and perchlorate in sodium hypochlorite solution used in water supply". 71 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 72 Adam, L. C., Fabian, I., Suzuki, K. and Gordon, G. "Hypochlorous Acid Decomposition in the pH 5-8 Region" Inorg. Chem. 1992, 31, 3534-3541. 73 Gordon, G., Cooper, W. J., Rice, R. G. and Pacey, G. E. Disinfectant Residual Measurement Methods, Second ed.; AWWA Research Foundation and American Water Works Association: Denver, 1992.

37

However, in general the rates of hypochlorite ion decomposition and chlorate ion

formation are dependent on the initial concentration of hypochlorite and chlorate ions,

temperature, pH, and ionic strength.74

Thus in order to begin to understand the factors impacting formation of

perchlorate ion in hypochlorite ion solutions, the following hypotheses were made:

perchlorate ion continuously forms in concentrated hypochlorite ion solutions, and the

rate of perchlorate ion formation is dependent on concentration of chlorate and

hypochlorite ions.

In addition, factors that impact decomposition of hypochlorite ion were also

studied. Transition metal ions have been shown to affect decomposition of hypochlorite

ion,75 and Cu2+ and Ni2+ were specifically shown to catalyze decomposition.76

Furthermore, iridium ion (Ir3+) is also known to catalyze decomposition of hypochlorite

ion,77 and as such, the effects of the following metal ions were investigated: transition

metal ions (Co2+, Cu2+, Fe3+, Mn2+, and Ni2+) and noble metal ions (Ag+, Au+, Ir3+ Pd2+,

and Pt2+).

Chlorite ion is produced during hypochlorite ion decomposition,78 shown by

Equation 17; however, chlorite ion is consumed by the reaction with hypochlorite ion to

produce chlorate ion, as shown by Equation 18. Because the rate of reaction, shown by

Equation 18, is faster, chlorite ion concentration reaches steady state, and typically is

present in trace amounts (typically 0.5 % of hypochlorite ion concentration).79

OCl- + OCl- → ClO2- + Cl- kClO2-(slow) (17)

OCl- + ClO2-→ ClO3

- + Cl- kClO3- (fast) (18)

Thus, if chlorite ion is added to a solution containing hypochlorite ion, it is

expected that the chlorate ion concentration will increase, as hypochlorite and chlorite

74 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 75 Lister, M. W. "Decomposition of Sodium Hypochlorite: The Catalyzed Reaction" Can. J. Chem. 1956, 34, 479–488. 76 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 77 Ayers, G. H. B., M. H. "Catalytic Decomposition of Hypochlorite Solution by Iridium Compounds. II. Kinetic Studies" J. Am. Chem. Soc. 1955, 77, 828-833. 78 Lister, M. W. "Decomposition of Sodium Hypochlorite: The Uncatalyzed Reaction" Can. J. Chem. 1956, 34, 465-478. 79 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion".

38

ions react. Based on these equations it was hypothesized that chlorite ion may also

potentially affect the formation of perchlorate ion.

Other potentially relevant species such as bromate ion (BrO3-), a known

contaminant found in sodium hypochlorite solutions,80 and bromide ion (Br-) were added

into the experimental matrix to determine any possible effects on the rate of perchlorate

ion formation.

3.1 Experimental Matrix and Chemicals

In this preliminary set of experiments to identify the major factors that affect the

rate of perchlorate ion formation, hypochlorite ion solutions at various concentrations of

hypochlorite and chlorate ions were prepared. In addition, transition metal ions, noble

metal ions, and bromide and bromate ions were added to hypochlorite ion solutions.

Spiked and non-spiked hypochlorite ion solutions were stored to allow decomposition of

hypochlorite ion to occur. As an aid in planning the incubation experiments, Bleach 2001

was used to predict decomposition of hypochlorite ion and formation of chlorate ion at

various concentrations of hypochlorite and chlorate ions, temperature, specific gravity,

pH (11-14), and sodium chloride concentrations. Because the decomposition of pH 11-

13 hypochlorite ion solutions at room temperature is relatively slow (the half life of 13 %

FAC solution at 25 ºC is 130 days), the preliminary experiments were performed at

elevated temperatures, so that multiple factors impacting perchlorate ion formation could

be screened over shorter time periods. Thus, a series of preliminary experiments were

performed at 75°, 60 ºC, and 50 ºC.

Bulk sodium hypochlorite solutions, 10–13% FAC, were obtained from Acros

Organics USA (Morris Plains, NJ) and VWR (Brisbane, CA). OSG sodium hypochlorite

solution, ~0.8 % FAC, was obtained from sodium hypochlorite on-site generator at River

Mountains Water Treatment Facility (Henderson, NV). Pure sodium chlorate (ACS

Grade, ≥99%) was obtained from VWR (Brisbane, CA) and added to hypochlorite ion

solutions, to generate solutions at various initial concentrations of chlorate ion. To

investigate the effects of chlorite ion, unstabilized sodium chlorite (Technical Grade,

80 Chlorine Institute, "Bromate in Sodium Hypochlorite Potable Water Treatment"

39

80%) was acquired from Acros Organics USA (Morris Plains, NJ) and added to sodium

hypochlorite solutions. To investigate the effects of transition metal ions, aliquots of

1,000 ppm Co2+, Cu2+, Fe3+, Ni2+, and Mn2+ standards (SPEX CertiPrep®, Inc.,

Metuchen, NJ) were spiked into sodium hypochlorite solutions. To investigate the effects

of noble metal ions on the rate of perchlorate ion formation, aliquots of 1,000 ppm Ag+,

Au+, Ir3+, Pd2+, and Pt2+ standards (Elements Inc., Shasta Lake, CA) were spiked to

sodium hypochlorite solutions. To determine effects of bromide and bromate ions, ACS

grade potassium bromide (Fisher Scientific, Pittsburgh, PA) and pure sodium bromate,

(99.5% min, EMD Chemicals Inc., Gibbstown, NJ) were spiked into sodium hypochlorite

solutions. Reagent Grade sodium chloride, (>99%, VWR, Brisbane, CA) was used to

spike sodium hypochlorite solutions as a proxy to increase the ionic strength.

Hypochlorite ion sample solutions were stored in 125 mL high density

polyethylene bottles. Analog and digital, general-purpose, heated water baths (VWR,

Brisbane, CA) were used for sample incubations. Temperatures were monitored daily,

using a glass laboratory mercury thermometer or a thermocouple thermometer with an

LCD screen (Fisher Scientific, Pittsburgh, PA). Sample aliquots were collected into 8

mL acid-washed, amber glass vials, cooled, and stored at 4 ºC prior to analysis. Samples

were analyzed by the potentiometric and iodometric titrations methods and by the LC-

MS/MS method, as described in Chapter 2.

3.2 Effect of Hypochlorite Ion Concentration

An aliquot of a stock, concentrated sodium hypochlorite solution was

analyzed by the potentiometric titration with sulfite ion, followed by the sequential

determination of chlorite and chlorate ions by the iodometric titration. Separate aliquots

of the stock hypochlorite ion solution were added to 100 mL volumetric flasks.

Calculations, based on the measured hypochlorite and chlorate ion concentrations of the

stock solution, were performed to determine the appropriate amount of sodium chlorate

to add to the individual diluted hypochlorite ion solutions. As a result the stock

hypochlorite ion solution and diluted hypochlorite ion solutions had the same

concentration of chlorate ion. Changes in the concentration of hypochlorite, chlorite,

40

chlorate, and perchlorate ions for duplicate samples were monitored during ten days of

incubation at 75 ºC in a water bath. Figure 10 shows overlaid smoothed-line plots of

hypochlorite ion decomposition and chlorate ion formation. Note: the downward sloping

smoothed lines represent hypochlorite ion decomposition, and the upward-sloping curves

show formation of chlorate ion shown in Figure 10.

As expected, solutions with the higher hypochlorite ion concentrations decompose

more rapidly, and form higher concentrations of chlorate ion, as depicted in Figure 10.

Figure 11 shows a smoothed-line plot of perchlorate ion formation for this set of samples.

Figure 10. Decomposition of hypochlorite ion and formation of chlorate ion at 75

ºC in solutions, at various initial concentrations of hypochlorite ion

0

200

400

600

800

1000

1200

0 2 4 6 8 10

Days

ClO

4-m

g/L

Form

ed

74g/L OCl- + 50 g/L ClO3-



Figure 11. Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions, at

various initial concentrations of hypochlorite ion

0

10

20

30

40

50

60

70

80

0 2 4 6 8 100

20

40

60

80

100

120

Days

OC

l-g/

L D

ecom

pose

d

ClO

3-g/

L Fo

rmed




41

A sample, with an initial concentration of 74 g/L OCl- produced significantly

higher concentration of perchlorate ion than samples with 50 g/L and 10 g/L OCl- (Figure

11). This supports the primary hypothesis that the rate of perchlorate ion formation is

strongly dependent on the concentration of hypochlorite ion.

3.3 Effect of Chlorate Ion Concentration

To investigate the effect of chlorate ion concentration on the rate of perchlorate

ion formation, a stock hypochlorite ion solution was divided into 100 mL aliquots. Pure

sodium chlorate (ACS Grade, ≥99%) was added to sample solutions to generate

hypochlorite solutions at various concentrations of chlorate ion. The sample solutions

were prepared and measured in duplicate. Measured concentrations of the hypochlorite,

chlorate, and perchlorate ions are reported as an average of the duplicate samples, and the

error is calculated as the difference between the average and individual duplicate

measurements. The addition of sodium chlorate caused the density of solutions to

increase and slight dilution of hypochlorite ion concentration was observed in some

samples. For example, the solution at 75 g/L OCl- decreased by 7% in hypochlorite ion

concentration, after a chlorate ion spike of 150 g/L was added. Figure 12 shows

smoothed-line plots of hypochlorite ion decomposition and formation of chlorate ion.

Similar to Figure 10, decomposition of hypochlorite ion at 75 ºC is rapid.

Figure 13 shows a smoothed-line plot of perchlorate ion formation as a function

of time. More perchlorate ion was formed in samples containing higher concentrations of

chlorate ion (Figure 13). This in turn supports the hypothesis that the rate of perchlorate

ion formation is dependent on concentration of chlorate ion.

42

0

10

20

30

40

50

60

70

80

0 2 4 6 8 10Days

OC

l-g/

L D

ecom

pose

d

020406080100120140160180200220

ClO

3-g/

L Fo

rmed

75 g/L OCl- + 24 g/L ClO3-

74 g/L OCl- + 74 g/L ClO3-

72 g/L OCl- + 124 g/L ClO3-

70 g/L OCl- + 174 g/L ClO3-


ºC in solutions, at various initial concentrations of chlorate ion

0

250

500

750

1000

1250

1500

1750

2000

2250

0 2 4 6 8 10Days

ClO

4-m

g/L

Form

ed

75 g/L OCl- + 24 g/L ClO3-

74 g/L OCl- + 74 g/L ClO3-

72 g/L OCl- + 124 g/L ClO3-

70 g/L OCl- + 174 g/L ClO3-

Figure 13. Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions, at

various initial concentrations of chlorate ion

43

3.4 Effect of Transition Metal Ions (Co2+, Cu2+, Fe3+, Mn2+, and Ni2+)

To investigate the effects of transition metal ions, aliquots of 1,000 ppm Co2+,

Cu2+, Fe3+, Ni2+, and Mn2+ standards (SPEX CertiPrep®, Inc., Metuchen, NJ) were

spiked into sodium hypochlorite solutions to achieve final concentration of 20 mg/L, 2

mg/L, and 0.2 mg/L of each metal ion. Following addition of a 20 mg/L spike, it was

observed that the decomposition of hypochlorite ion was very rapid, and no change in

concentration of perchlorate ion was observed after 30 days of incubation at 75 ºC. Even

at a 2 mg/L spike, the catalyzed decomposition of hypochlorite ion is still rapid, as

compared to the control sample shown in Figure 14. The decomposition of hypochlorite

ion is clearly catalyzed by the presence of Co2+, Cu2+, Fe3+, Mn2+, and Ni2+, and the rate

of decomposition increases with higher concentrations of these cations. At the same time

less perchlorate ion was formed as a function of time in metal ion spiked samples, as

shown in Figure 15.

0

10

20

30

40

50

60

70

80

0 1 2 3 4 5 6 7 8 9 10Days

0

10

20

30

40

50

60

70

8075 g/L OCl- + 24 g/L ClO3- + 2 mg/L Me

75 g/L OCl- + 24 g/L ClO3- + 0.2 mg/L Me

OC

l-g/

L D

ecom

pose

d

ClO

3-g/

L Fo

rmed

75 g/L OCl- + 24 g/L ClO3- (Control)

Figure 14. Effects of Transition Metals Ions (Me = Co2+, Cu2+, Fe3+, Mn2+, and

Ni2+) on the hypochlorite ion decomposition and the chlorate ion formation

Observations, drawn from Figures 14 and 15, suggest that the presence of Co2+,

Cu2+, Fe3+, Mn2+, and Ni2+ even at a concentration of 0.2 mg/L exhibits a much stronger

effect on the decomposition of hypochlorite ion such that no additional perchlorate ion

44

formation was observed. Thus, these metal ions can not enhance the rate of perchlorate

ion formation under these conditions (pH 11-13, FAC > 5.0 %), because the primary

reactant to produce perchlorate ion (i.e. hypochlorite ion) is decomposed rapidly.

0

100

200

300

400

500

600

700

0 1 2 3 4 5 6 7 8 9 10Days

75 g/L OCl- + 24 g/L ClO3- (Control)

75 g/L OCl- + 24 g/L ClO3- + 2 mg/L Me

75 g/L OCl- + 24 g/L ClO3- + 0.2 mg/L Me

ClO

4-m

g/L

Form

ed

Figure 15. Effects of Transition Metals Ions (Me = Co2+, Cu2+, Fe3+, Mn2+, and

Ni2+) on perchlorate ion formation

3.5 Effect of Noble Metal Ions (Ag+, Au+, Ir3+, Pd2+, and Pt2+)

Similar to the investigation of transition metal ion effects on the rate of

perchlorate ion formation, aliquots of 1,000 ppm Ag+, Au+, Ir3+, Pd2+, and Pt2+ standards

(Elements Inc., Shasta Lake, CA) were spiked into sodium hypochlorite solutions. The

final concentration of each metal ion was 0.2 mg/L. Spiked samples were divided into

two groups: a set with a spike of 0.2 mg/L Ag+, Au+, Ir3+, Pd2+, and Pt2+, and a set with a

spike of 0.2 mg/L Ir3+ only. This was done to separate the effects of Ir3+, which has been

linked to enhanced decomposition of hypochlorite ion,81 and other metals in this group.

Finally, spiked samples and the control sample (no spike) were incubated at 50 ºC, and

the target analytes were monitored. Figure 16 shows overlaid smoothed-line plots of

hypochlorite ion decomposition and formation of chlorate ion.

81 Ayers "Catalytic Decomposition of Hypochlorite Solution by Iridium Compounds. II. Kinetic Studies".

45

01020304050607080

0 1 2 3 4 5 6Days

0

10

20

30

40

50

60Set 1: 0.2 mg/L Noble Me IonsSet 2: 0.2 mg/L Iridium Ion

No Spike (Control: 73 g/L OCl- + 27 g/L ClO3-)

OC

l-g/

L D

ecom

pose

d

ClO

3-g/

L Fo

rmed

Figure 16. Effects of noble metal ions (Noble Me = Ag+, Au+, Ir3+

, Pd2+, and Pt2+), plots of hypochlorite ion decomposition and chlorite ion formation

It is evident from Figure 16 that at 0.2 mg/L spike, unlike the transition metal ions, the

noble metal ions (Ag+, Au+, Ir3+, Pd2+, and Pt2+) do not enhance decomposition of

hypochlorite. Thus, it was no surprise when no significant differences were observed in

the concentration of perchlorate ion formed as a function of time in these samples, as

shown in Figure 17.

153045607590

105120135150165

0 1 2 3 4 5 6 7 8 9 10Days

ClO

4-m

g/L

Form

ed

Set 1: 0.2 mg/L Noble Me IonsSet 2: 0.2 mg/L Iridium Ion


Figure 17. Effects of noble metals ions (Noble Me = Ag+, Au+, Ir3+

, Pd2+, and Pt2+), overlaid plots of perchlorate ion formation

46

Table 14 shows the changes in concentration of perchlorate ion over time of

spiked samples and the control samples. This illustrates that the statistical difference in

results is relatively small, with average relative standard deviation (RSD) of 3.1 % ± 1.2.

The average standard deviation (SD) of all the samples is quite comparable to the

difference of duplicate samples, such that differences between spiked and control sample

solutions are statistically insignificant. Furthermore these noble metal ions have not been

reported to be typically present in hypochlorite ion solutions above 0.2 mg/L, and thus

are not potentially significant species involved in formation of perchlorate ion in

hypochlorite ion solutions.

Table 14. Changes in perchlorate ion concentration of samples spiked with Ag+, Au+, Ir3+

, Pd2+, and Pt2+ (Noble Me) vs. control (no spike), incubated at 50 ºC

ClO4- (mg/L)

Day Sample Index: 0 1 2 3 6 8 10

Set 1: 0.2 mg/L Noble Me 20.4 37.6 50.5 71.5 108 135 161 Set 1: 0.2 mg/L Noble Me (Duplicate) 20.7 37.1 52.0 70.9 112 137 150 Set 2: 0.2 mg/L Ir 19.7 40.3 54.0 69.5 115 128 161 Set 2: 0.2 mg/L Ir (Duplicate) 20.2 40.5 52.0 72.5 115 140 162 No spike: Control 20.0 34.6 52.0 74.4 114 138 157 No spike: Control (Duplicate) 19.9 37.8 50.0 69.4 115 139 161

Mean 20.1 38.0 52.0 71.4 113 136 159 Std. Dev. 0.4 2.2 1.4 1.9 2.8 4.4 4.6 RSD (%) 2.0 5.8 2.7 2.7 2.5 3.2 2.9

3.6 Effect of Chlorite Ion Concentration

To investigate the effects of chlorite ion concentration, unstabilized sodium

chlorite (Technical Grade, 80%) was added to sodium hypochlorite solutions to achieve a

ClO2- spike at 15 g/L. In addition to the chlorite ion spike, separate hypochlorite ion

solutions were spiked with pure sodium chlorate (ACS Grade, ≥99%) at 15 g/L ClO2-

and/or at 100 g/L ClO3- to determine any combined effects of the two anions. Sample

solutions prepared in duplicate were incubated at 50 ºC in a water bath. In addition, a

similar sample set was incubated at 30 ºC and used to confirm any observable effects of

47

chlorite ion concentration on the rate of perchlorate ion formation over prolonged periods

of time.

Figure 18 shows overlaid plots of hypochlorite ion decomposition and chlorate

ion formation at 50 ºC and 30 ºC. Since changes in concentration of hypochlorite and

chlorate ions are expected due to addition of chlorite ion, samples were labeled using

molar product, which was calculated by multiplying concentrations of hypochlorite and

chlorate ions, shown in Equation 19.

][][ 3−− ×= ClOOClMP (19)

0

20

40

60

80

100

0 1 2 3 4 5 6 7 8 9 10Days

020406080100120140160

0.26 M MP ( Control )0.26 M MP + 15 g/L ClO2

-

2.09 M MP + 15 g/L ClO2- + 100 g/L ClO3

-

2.13 M MP + 100 g/L ClO3-

OC

l-g/

L D

ecom

pose

d

ClO

3-g/

L Fo

rmed

0

20

40

60

80

100

0 25 50 75 100 125 150 175 200Days

020406080100120140160

OC

l-g/

L D

ecom

pose

d


-

2.07 M MP + 15 g/L ClO2- +100 g/L ClO3

-

2.09 M MP + 100 g/L ClO3-

ClO

3-g/

L Fo

rmed

Figure 18. Plots of hypochlorite ion decomposition and chlorate ion formation, in

solutions at various initial concentrations of chlorite ion and/or chlorate ion at (a) 50 ºC, (b) 30 ºC

(a)

(b)

48

As expected, hypochlorite ion solutions at 15 g/L ClO2- spike had noticeably more

chlorate ion formed than the control solution, while also showing faster hypochlorite ion

decomposition, as shown in Figure 18 (a) and (b). This is due to reaction between

hypochlorite and chlorite ions, shown in Equation 20:

OCl- + ClO2-→ ClO3

- + Cl- kClO3- (fast) (20)

If the concentration of chlorate ion is increased, additional amounts of perchlorate

ion can be expected to form. However, as was shown earlier, the formation of

perchlorate ion is dependent on both the concentration of chlorate and hypochlorite ions,

and so the effects of chlorite ion are two-fold. The first effect is due to the conversion of

chlorite ion to chlorate ion, enhancing the formation of perchlorate ion. The second

effect, occurring at the same time, is reduction of hypochlorite ion concentration, which

would decrease the rate of perchlorate ion formation. Overlaid smoothed-line plots of

perchlorate ion formation at 50° and 30 ºC are shown in Figure 19 (a); (b).

For further analysis, changes in molar product of the samples spiked with chlorite

ion and chlorate ion versus just chlorate ion were plotted and overlaid with formation of

perchlorate ion, shown in Figure 20. The observed similarities in changes of the molar

product of the chlorate ion-spiked sample and chlorite/chlorate ion-spiked sample, and

similarities observed in the formation of perchlorate ion indicate that the addition of the

chlorite ion does not appear to offer a substantial change in perchlorate ion formation

beyond what would be observed from addition of the chlorate ion.

49

0

100

200

300

400

500

600

0 1 2 3 4 5 6 7 8 9 10Days

ClO

4-m

g/L

Form

ed


-

2.09 M MP + 15 g/L ClO2- + 100 g/L ClO3

-

2.13 M MP + 100 g/L ClO3-

0

100

200

300

400

500

600

0 25 50 75 100 125 150 175 200Days


-

2.07 M MP + 15 g/L ClO2- + 100 g/L ClO3

-

2.09 M MP + 100 g/L ClO3-

ClO

4-m

g/L

Form

ed

Figure 19. Overlaid plot of changes in molar product and perchlorate ion

formation over time in solutions at various initial concentrations of chlorite ion and/or chlorate ion at (a) 50 ºC, (b) 30 ºC

Thus, it was concluded that although the addition of chlorite ion to hypochlorite

ion solution did affect the formation of perchlorate ion, the effect was based on reaction

with hypochlorite ion to produce chlorate ion. Furthermore, the addition of chlorite and

chlorate ions together did not show additional amounts of perchlorate ion formed as

shown in Figures 19 and 20. Therefore, chlorite ion is not directly involved in the

formation of the perchlorate ion.

(a)

(b)

50

0.500.700.901.101.301.501.701.902.102.30

0 1 2 3 4 5 6 7 8 9 10Days

Mol

ar P

rodu

ct

075150225300375450525600

2.13 M MP + 100 g/L ClO3-

2.09 M MP + 15 g/L ClO2- + 100 g/L ClO3

-

ClO

4-m

g/L

Form

ed

0.500.700.901.101.301.501.701.902.102.30

0 25 50 75 100 125 150 175 200

Days

Mol

ar P

rodu

ct

075150225300375450525600

2.09M MP + 100 g/L ClO3-

2.07M MP + 15 g/L ClO2- + 100 g/L ClO3

-

ClO

4-m

g/L

Form

ed

Figure 20. Plots of perchlorate ion formation in solutions at various initial

concentrations of chlorite ion and/or chlorate ion, (a) 30 ºC, (b) 50 ºC

3.6.1 Combined Effect of Transition Metal Ions, Chlorite and Bromide Ions

To determine if there are potentially multiple pathways to form perchlorate ion,

the combined effects of chlorite ion, transition metals ions (Co2+, Cu2+, Fe3+, Mn2+, and

Ni2+), and bromide ion were investigated. Samples were grouped into the following sets:

Set 1: ClO2- at 15 g/L + Transition Metals Spike at 0.2 mg/L

Set 2: ClO2- at 15 g/L + ClO3

- at 100 g/L + Transition Metals Spike at 0.2 mg/L

Set 3: ClO2- at 15 g/L + Br- at 15 g/L

(a)

(b)

51

Pure potassium bromide (ACS Grade) was used for the preparation of bromide

ion spiked solutions. All sample solutions were prepared in duplicate, and incubated at

50 ºC. The reaction between hypochlorite ion and bromide ion, in the pH 10-14 region,82

can be described by Equation 21.

OCl- + Br- → OBr- + Cl- (21)

Thus, similar to chlorite ion, bromide ion is expected to readily react with

hypochlorite ion. The produced hypobromite ion (OBr-) decomposes to produce bromite

and bromate ions,83 as shown in Equations 22 and 23.

OBr- + OBr- → BrO2- + Br- (22)

BrO2- + OBr- → BrO3

- + Br- (23)

Figure 21 shows overlaid smoothed-line plots of hypochlorite ion decomposition

and chlorate ion formation. As, expected, sample solutions spiked with transition metal

ions decompose rapidly, due to the catalytic effects of these ions. Sample solution spiked

with bromide ion results in rapid reaction between hypochlorite and bromide ions, as

described by Equation 21. However, once this reaction is completed (by the first day of

incubation), the rapid decomposition of hypochlorite ion slows down, and for the

remainder of experiment, normal decomposition of hypochlorite ion is observed. Figure

22 shows smoothed-line plots of perchlorate ion formation.

82 Farkas, L., Lewin, M. and Bloch, R. "The Reaction between Hypochlorite and Bromides" J. Am. Chem. Soc. 1949, 71, 1988-1991. 83 Perlmutter-Hayman, B. and Stein, G. "The Kinetics of the Decomposition of Alkaline Solutions of Hypobromite-Specific Ionic Effects on Reaction Rate" J. Phys. Chem. 1959, 63, 734-738.

52

01020304050607080

0 1 2 3 4 5 6Days

OC

l-g/

L D

ecom

pose

d

Set 1: 15 g/L ClO2- + 0.2 mg/L Me

Set 3: 15 g/L ClO2- + 15 g/L Br-


Set 2: 15 g/L ClO2- + 100 g/L ClO3

- + 0.2 mg/L Me


ºC in solutions spiked with bromide, chlorite, and transition metal ions (Me = Co2+, Cu2+, Fe3+, Mn2+, and Ni2+)

0

50

100

150

200

0 1 2 3 4 5 6 7 8 9 10Days

ClO

4-m

g/L

Form

ed

Set 1: 15 g/L ClO2- + 0.2 mg/L Me

Set 2: 15 g/L ClO2- + 100 g/L ClO3

- + 0.2 mg/L MeSet 3: 15 g/L ClO2

- + 15 g/L Br-


Figure 22. Formation of perchlorate ion at 50 ºC in solutions spiked with

bromide, chlorite, and transition metal ions (Me = Co2+, Cu2+, Fe3+, Mn2+, and Ni2+)

Figure 22 shows that the non-spiked hypochlorite ion solution (control) produced

considerably more perchlorate ion than in solutions of Set 1 and 3. Thus, it can be

concluded that chlorite ion and transition metal ions, as well as the chlorite and bromide

ion combination do not enhance the formation of perchlorate ion. The solution, spiked

with chlorite/chlorate ions and transition metal ions, produced higher amounts of

53

perchlorate ion than the control solution, due to the higher concentration of chlorate ion

than in the control solution. However, because hypochlorite ion was catalytically-

decomposed, the rate of perchlorate ion formation decreased. Thus, based on the

investigation of the role of chlorite ion, it can be concluded that, there is no additional

pathway to form perchlorate ion, when chlorite ion, bromide ion, and/or transition metals

ions are present. Therefore, these species are not directly involved in formation of

perchlorate ion.

3.7 Effect of Bromide Ion and Bromate Ion Concentration

Conversion of bromide ion to bromate ion was expected by a pathway given by

equations 22-24:

OCl- + Br- → OBr- + Cl- (21)

OBr- + OBr- → BrO2- + Br- (22)

BrO2- + OBr- → BrO3

- + Br- (23)

Thus, in theory the presence of an additional oxidizing agent, such as

hypobromite ion and/or bromite ion, may potentially alter the perchlorate ion formation

pathway. To determine the effects of bromide and bromate ions, potassium bromide,

ACS grade (Fisher Scientific, Pittsburgh, PA), and sodium bromate, 99.5% min (EMD

Chemicals Inc., Gibbstown, NJ), were added to the sodium hypochlorite solutions as

follows:

Set 1: Br- spike at 15 g/L

Set 2: BrO3- spike at 15 g/L

Set 3: Br- + BrO3- spike at 15 g/L

Set 4: Br- spike at 30 g/L

Set 5: BrO3- spike at 30 g/L

Set 6: Br- + BrO3- spike at 30 g/L

Figure 23 shows overlaid plots of hypochlorite ion decomposition in various

hypochlorite ion solutions at pH 12.5. As expected, samples spiked with bromide ion

show a loss in hypochlorite ion initially, and then decompose normally. Figure 23 shows

overlaid plots of bromate ion formation.

54

0102030405060708090

0 1 2 3 4 5 6Days

Set 1: 15 g/L Br-

Set 2: 15 g/L BrO3-

Set 3: 15 g/L Br- + BrO3-


OC

l-g/

L D

ecom

pose

d

Set 4: 30 g/L Br-

Set 5: 30 g/L BrO3-


Figure 23. Decomposition of hypochlorite ion at 50 ºC in solutions spiked with

bromide and bromate ions at pH~12.5

0

20

40

60

80

100

0 1 2 3 4 5 6 7 8 9 10Days

BrO

3-g/

L Fo

rmed

No Spike (BrO3- at 70 mg/L)

Set 1: 15 g/L Br-

Set 2: 15 g/L BrO3-


Set 4: 30 g/L Br-

Set 5: 30 g/L BrO3-



and bromate ions at pH~12.5 Bromide spiked samples show rapid formation of bromate ion, as can be seen in

Figure 24, which coincides with rapid decomposition of hypochlorite ion that can be

observed in Figure 23. This suggests that if any hypobromite ion is formed, it rapidly

decomposes to bromate ion, similar to the decomposition pathway of the hypochlorite

ion. Additionally, no change in bromate ion concentration or any enhanced

decomposition of hypochlorite ion was observed in samples that were spiked with only

55

bromate ion. Furthermore, Figure 25 show that perchlorate ion formation in bromate ion

spiked samples was identical to control sample (non-spiked sample). This would strongly

suggest that bromate ion is a spectator species that is not involved in formation of

perchlorate ion or decomposition of hypochlorite ion.

Addition of bromide ion caused a decrease in hypochlorite ion concentration,

which in turn lowered the amount of perchlorate ion formed. Thus, similar to effects of

transition metal ions, bromide ion is not involved in perchlorate ion formation; however,

it may affect the rate of perchlorate ion formation due to its reaction with hypochlorite

ion.

020406080

100120140160

0 1 2 3 4 5 6 7 8 9 10Days


ClO

4-m

g/L

Form

ed

Set 1: 15 g/L Br-

Set 2: 15 g/L BrO3-


Set 4: 30 g/L Br-

Set 5: 30 g/L BrO3-



and/or bromate ions at pH~12.5

3.8 Effect of Ionic Strength

The rate of hypochlorite ion decomposition and chlorate ion formation increases

with an increase in the ionic strength, and these effects have been thoroughly quantified

previously.84 As a result it was shown that the rate constant for the hypochlorite ion

84 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

56

decomposition depends on the ionic strength and pH.85 Ionic strength (μ), defined as the

half sum of all the ions, (Equation 24), can be increased with addition of an electrolyte,

221

ii

i zc∑=μ (24)

where, μ is ionic strength (mol/L), ci concentration (mol/L) of the ith species, and zi is its

charge.86 For illustration purposes, Bleach 2001, was used to predict decomposition of

hypochlorite ion at different initial concentrations of chloride ion, shown in Figure 26.

Similarly, formation of chlorate ion was plotted as a function of initial chloride ion

concentration, as shown in Figure 27. The following parameters were used in calculations

by Bleach 2001: initial concentration of 80 g/L OCl-, 0.002 g/L ClO3-, pH = 12.5, and

temperature = 40 ºC.

0

10

20

30

40

50

60

70

80

90

0 7 14 21 28 35 42 49 56 63Days

OC

l-g/

L D

ecom

pose

d 80 g/L OCl- + 35.5 g/L Cl-80 g/L OCl- + 106 g/L Cl-80 g/L OCl- + 177 g/L Cl-

Figure 26. Decomposition of hypochlorite ion at 40 ºC in solutions at various

initial concentrations of chloride ion

Pure sodium chloride (Reagent Grade, >99%) was used to spike sodium

hypochlorite solutions as a proxy to increase the ionic strength. Samples were incubated

at 40 ºC.

85 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 86 Harris, D. C. Quantitative Chemical Analysis, 5th ed.; W.H. Freeman and Company: New York, 2001.

57

0

5

10

15

20

25

30

35

40

0 7 14 21 28 35 42 49 56 63Days

ClO

3-g/

L Fo

rmed

80 g/L OCl- + 35.5 g/L Cl-80 g/L OCl- + 106 g/L Cl-80 g/L OCl- + 177 g/L Cl-

Figure 27. Formation of chlorate ion at 40 ºC in solutions at various initial concentrations of chloride ion

Figure 28 shows overlaid plots of perchlorate ion formation in sodium

hypochlorite solutions with variable concentrations of chloride ion. Clearly, more

perchlorate ion was formed in samples containing higher concentration of chloride ion.

Thus, it is concluded that the rate of perchlorate ion formation increases at higher ionic

strength. In the bulk, concentrated sodium hypochlorite solutions, concentration of

chloride ion equaled or was greater than hypochlorite ion. In general significant

differences in the ionic strength of the bulk sodium hypochlorite solution were not

expected.

0255075

100125150

0 7 14 21 28 35 42 49 56 63 70 77Day

No Spike (Control: 0.48 M OCl- + 0.62 M ClO3-)

Cl- Spike at 25 g/LCl- Spike at 50 g/LCl- Spike at 100 g/L

ClO

4-m

g/L

Form

ed

Figure 28. Formation of perchlorate ion at 40 ºC in solutions at various initial

concentrations of chloride ion

58

3.9 Effect of pH

Previously, it was identified that the rate of hypochlorite ion decomposition

depends on pH and is at a minimum in the pH 11-13 range.87 In hypochlorite ion

solutions at pH above 13, the hypochlorite ion decomposition is enhanced due to the high

concentration of hydroxide ion, which increases the ionic strength. Below pH 11, the

acid-catalyzed decomposition of hypochlorite ion begins to occur, such that below pH 9

the decomposition of hypochlorite becomes third-order in hypochlorite.88 Because

adjusting the pH will also change the ionic strength, it was hypothesized that the rate of

perchlorate ion formation may be affected by pH.

To investigate the effect of lower pH values, aliquots of 1.4 M and 0.9 M OCl-

solutions were adjusted with hydrochloric acid and incubated at 40 ºC. Figure 29 shows

overlaid decomposition plots of the hypochlorite ion and formation plots of the chlorate

ion. Figure 30 shows overlaid formation plots of the perchlorate ion.

Hypochlorite ion solutions at pH 11, have faster decomposition of hypochlorite

ion than at pH 13, and this resulted in lower amount of perchlorate ion formed at pH 11,

as can be seen in Figure 30. At pH 9, the decomposition of hypochlorite ion was too rapid

(in a matter of hours 1.4 M OCl- decomposed to < 0.1 M OCl-, Figure 29) and the effect

of pH 9 during this experiment on the rate of perchlorate ion formation was inconclusive.

87 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 88 Adam, Fabian, Suzuki and Gordon "Hypochlorous Acid Decomposition in the pH 5-8 Region".

59

00.20.40.60.81.01.21.4

0 7 14 21 28 35Days

[OC

l- ]

0

0.2

0.4

0.6

0.8

1.0

[ClO

3- ]

1.4 M OCl- + 0.6 M ClO3- pH 13

1.4 M OCl- + 0.6 M ClO3- pH 11

1.4 M OCl- + 0.6 M ClO3- pH 9

00.10.20.30.40.50.60.70.80.91.0

0 7 14 21 28 35Days

[OC

l-]

00.10.20.30.40.50.60.70.80.91.0

[ClO

3-]

0.9 M OCl- + 0.6 M ClO3- pH 13

0.9 M OCl- + 0.6 M ClO3- pH 11

0.9 M OCl- + 0.6 M ClO3- pH 9


ºC in (a) 1.4 M OCl-, (b) 0.9 M OCl- solutions at various initial pH

(a)

(b)

60

050

100150200250300350400

0 7 14 21 28 35 42 49 56 63 70 77Days

1.4 M OCl- + 0.6 M ClO3- pH 13

1.4 M OCl- + 0.6 M ClO3- pH 11

1.4 M OCl- + 0.6 M ClO3- pH 9

ClO

4-m

g/L

Form

ed

0255075

100125150175200

0 7 14 21 28 35 42 49 56 63 70 77Days

ClO

4-m

g/L

Form

ed

0.9 M OCl- + 0.6 M ClO3- pH 13

0.9 M OCl- + 0.6 M ClO3- pH 11

0.9 M OCl- + 0.6 M ClO3- pH 9

Figure 30. Formation of perchlorate ion at 40 ºC in (a) 1.4 M OCl-, (b) 0.9 M

OCl- solutions at various initial pH

As a follow-up experiment to investigate the effects of pH in 9-11 region, aliquots

of an OSG sodium hypochlorite solution were adjusted to different pH values using

sodium hydroxide. Samples solutions at 0.12 M OCl- were adjusted to pH 9.35, 10.65,

11.9, and 13.3. Because decomposition of 0.12 M OCl- solution is quite slow, samples

were incubated at 60 ºC. Figure 31 shows overlaid decomposition plots of hypochlorite

ion and formation plots of chlorate ion for these samples.

(a)

(b)

61

01234567

0 5 10 15 20 25Days

0.12 M OCl- + 0.006 M ClO3- pH 9.35

0.12 M OCl- + 0.006 M ClO3- pH 10.65

0.12 M OCl- + 0.006 M ClO3- pH 11.90

0.12 M OCl- + 0.006 M ClO3- pH 13.30

OC

l-g/

L D

ecom

pose

d

0

0.5

1

1.5

2

2.5

3

0 5 10 15 20 25Days

0.12 M OCl- + 0.006 M ClO3- pH 9.35

0.12 M OCl- + 0.006 M ClO3- pH 10.65

0.12 M OCl- + 0.006 M ClO3- pH 11.90

0.12 M OCl- + 0.006 M ClO3- pH 13.30

ClO

3-g/

L Fo

rmed

Figure 31. Overlaid plots of (a) hypochlorite ion decomposition; (b) chlorate ion

formation at 60 ºC 0.12 M OCl- solutions at various initial pH

As expected, hypochlorite ion solution at pH 9.35 had the fastest rate of

hypochlorite ion decomposition, and the fastest rate of chlorate ion formation as shown in

Figures 31 (a) and (b). Interestingly, the rate of perchlorate ion formation was enhanced

in the sodium hypochlorite solution, having an initial pH of 9.35, as shown in Figure 32.

This would suggest that in dilute hypochlorite ion solutions (such as OSG sodium

hypochlorite), the rate of perchlorate ion formation may be also dependent on pH.

However, from a practical stand-point, perchlorate ion formation is more dependent on

the concentration of hypochlorite and chlorate ions. Thus the effect of pH is relatively

insignificant for bulk, concentrated sodium hypochlorite solutions.

(a)

(b)

62

0

100

200

300

400

500

600

0 3 6 9 12 15 18 21Days

0.12 M OCl- + 0.006 M ClO3- pH 9.35

0.12 M OCl- + 0.006 M ClO3- pH 10.65

0.12 M OCl- + 0.006 M ClO3- pH 11.90

0.12 M OCl- + 0.006 M ClO3- pH 13.30

ClO

4-μg

/L F

orm

ed

Figure 32. Overlaid plots of perchlorate ion formation at 60 ºC 0.12 M OCl-

solutions at various initial pH

3.10 Conclusions

The results of the preliminary experiments, presented in this chapter, indicate that

perchlorate ion formation in sodium hypochlorite solutions is dependent on several

factors: (1) Concentration of hypochlorite and chlorate ions directly impact perchlorate

ion formation; (2) Presence of transition metal ions, chlorite ion or bromide ion indirectly

decrease perchlorate ion formation, by reactions with hypochlorite ion; (3) Presence of

noble metal ions or bromate ion has no observable effect on perchlorate formation; (4)

An increase in chloride ion concentration, as a proxy to increase ionic strength, enhances

perchlorate ion formation and thus would need to be accounted for; (5) pH effects in

concentrated hypochlorite ion solutions are more dominant on hypochlorite ion

decomposition (faster kinetics) than on perchlorate ion formation; however, in more

dilute (i.e. more stable) hypochlorite ion solutions, perchlorate ion formation also appears

to be acid-catalyzed.

63

CHAPTER 4. KINETICS OF PERCHLORATE ION FORMATION AND

DETERMINATION OF THE RATE LAW

The objectives for this portion of the study were to elucidate the rate law for the

formation of perchlorate ion and determine the rate constant(s). The effects of

temperature and ionic strength on the rate constant(s) were investigated in order to

provide a readily usable model (with the fewest number of parameters) for perchlorate

ion formation.

As was shown in the previous chapter, both hypochlorite and chlorate ion

concentrations were found to have a strong effect on the rate of perchlorate ion

formation. An increase in concentration of hypochlorite or chlorate ions consistently

resulted in an increased rate of perchlorate ion formation and the final amount of

perchlorate ion formed. Therefore, the hypothesis that perchlorate ion formation was a

direct result of reactions between hypochlorite and chlorate ions was verified. The most

general stoichiometric reaction, given by Equation 25 was assumed.

−−−− +→+ ClClOClOOCl 43 (25)

Because perchlorate ion formation was dependent on both concentrations of

hypochlorite and chlorate ions, the rate law can be expressed by Equation 26.

pm ClOOClkdt

ClOdRate ][][][32

4 −−−

×== (26)

By taking the natural log of both sides of Equation 26, an expression shown in

Equation 27 is obtained:

]ln[]ln[ln)ln( 32−− ×+×+= ClOpOClmkRate (27)

The rate of perchlorate ion formation and the concentrations of hypochlorite and

chlorate ions can be measured experimentally. The second-order rate constant (k2) can be

determined, if the reaction orders with respect to hypochlorite ion and chlorate ion (m and

p) are known.

To determine the reaction order with respect to hypochlorite and chlorate ions, a

series of experiments were designed where the concentration of hypochlorite or chlorate

ions was varied while holding the concentration of the other reactant constant. Based on

64

the fact that the rate of hypochlorite ion decomposition and the rate of chlorate ion

formation are temperature-dependent,89 these experiments were performed at multiple

temperatures. To quantify only the effects of hypochlorite and chlorate ion concentration

on perchlorate ion formation, the experiments were conducted on hypochlorite ion

solutions with pH in the range of 12-13, where the acid-catalyzed decomposition of

hypochlorite ion is at minimum.90 This was done to avoid variation in the rate of

perchlorate ion formation due to significant changes in pH and ionic strength.

4.1 Reaction Order with Respect to Chlorate Ion: ln (d[ClO4-]/dt) vs. ln [ClO3

-]

In this set of experiments, hypochlorite ion solutions were prepared in duplicate

with at least 70 g/L OCl-, as the initial concentration, while adding chlorate ion at 50 g/L,

100 g/L, and 150 g/L. A control hypochlorite ion solution containing low chlorate ion

concentration was used. Samples were incubated at 30°, 40°, 50°, and 75 ºC.

Figure 33 (a) shows overlaid plots of perchlorate ion formation as a function of

initial chlorate ion concentration, and (b) the effects of initial chlorate ion concentrations

on the decomposition of hypochlorite ion and formation of chlorate ion at 30 ºC. Figures

34 and 35 show the effects of chlorate ion concentration during incubation experiments at

40 ºC and at 50 ºC. Note: the downward sloping smoothed lines represent hypochlorite

ion decomposition, and the upward-sloping curves show formation of chlorate ion in

Figures 33-35 (b).

As was observed during the preliminary experiment at 75 ºC reported in Chapter

3, Figure 13 page 42, the addition of chlorate ion to the hypochlorite ion solution results

in a proportionate increase in perchlorate ion. This confirms that the rate of perchlorate

ion formation changes as a function of initial concentration of chlorate ion. The initial

concentration of hypochlorite ion was kept constant. However, as was discussed in

Chapter 3, with addition of sodium chlorate, a small dilution of hypochlorite ion

concentration occurred in some samples. In general, these dilutions were less than 5% of

the non-spiked hypochlorite ion solutions, and thus the initial concentration of

89 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 90 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

65

hypochlorite ion was assumed to be constant. Table 15 shows the changes in

hypochlorite ion concentration in sodium hypochlorite solutions with various

concentrations of chlorate ion.

0100200300400500600700800

0 20 40 60 80 100 120 140 160 180 200Days

ClO

4-m

g/L

84 g/L OCl- + 15 g/L ClO3-84 g/L OCl- + 15 g/L ClO3-




0102030405060708090

100

0 25 50 75 100 125 150 175 200Days

OC

l-g/

L D

ecom

pose

d0

50

100

150

200

250

ClO

3-g/

L Fo

rmed

84 g/L OCl- + 15 g/L ClO3-

82 g/L OCl- + 63 g/L ClO3-

80 g/L OCl- + 112 g/L ClO3-

77 g/L OCl- + 159 g/L ClO3-

Figure 33. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of

hypochlorite ion and formation of chlorate ion at 30 ºC in solutions at various initial concentrations of chlorate ion at pH ~12.5

0100200300400500600700800

0 7 14 21 28 35 42 49 56 63 70 77Days

ClO

4-m

g/L





0

20

40

60

80

100

0 5 10 15 20 25 30 35Days

0

50

100

150

200

25070 g/L OCl- + 149 g/L ClO3

-70 g/L OCl- + 149 g/L ClO3-




OC

l-g/

L D

ecom

pose

d

ClO

3- g/L

For

med



(b) (a)

(b) (a)

66

0100200300400500600700800

0 2 4 6 8 10Days

ClO

4-m

g/L

85 g/L OCl- + 13 g/L ClO3-

84 g/L OCl- + 62 g/L ClO3-

83 g/L OCl- + 110 g/L ClO3-

80 g/L OCl- + 154 g/L ClO3-

0102030405060708090

100

0 2 4 6 8 10Days

OC

l-g/

L D

ecom

pose

d

0

50

100

150

200

250

ClO

3-g/

L Fo

rmed

85 g/L OCl- + 13 g/L ClO3-

84 g/L OCl- + 62 g/L ClO3-

83 g/L OCl- + 110 g/L ClO3-

80 g/L OCl- + 154 g/L ClO3-



Table 15. Decomposition of hypochlorite ion at 30 ºC in solutions, at various initial concentrations of chlorate ion ([ClO3

-]0) at pH ~12.5 [OCl-] (mol/L)

Days [ClO3

-]0 15 g/L

[ClO3-]0

63 g/L [ClO3

-]0 112 g/L

[ClO3-]0

159 g/L Mean

(mol/L)Std. Dev.

RSD (%)

0 1.637 1.594 1.557 1.506 1.574 0.056 3.5 11 1.423 1.372 1.322 1.269 1.347 0.066 4.9 21 1.241 1.200 1.148 1.092 1.170 0.065 5.5 34 1.108 1.062 1.012 0.958 1.035 0.065 6.2 49 0.953 0.917 0.877 0.799 0.887 0.066 7.5 63 0.858 0.837 0.777 0.730 0.801 0.058 7.3

As shown in Table 15, the difference in hypochlorite ion concentration becomes

more significant over longer incubation periods. The rates of perchlorate ion formation

were determined during the first thirty-four days of each chlorate ion spike experiment, in

order to minimize the error based on differences in concentration of hypochlorite ion

(after 34 days the difference in hypochlorite ion was 6.2 %, as shown in Table 15). To

calculate the rates of perchlorate ion formation, the change in measured concentration of

perchlorate ion was divided by the incubation interval. For example, after eleven days of

incubation, concentration of perchlorate ion in hypochlorite ion solution with [OCl-]0 =

1.637 M and [ClO3-]0 = 0.178 M increased from 7.10 mg/L ClO4

- to 12.45 mg/L ClO4-,

(b) (a)

67

based on the average of duplicate samples. Thus, 53.8 μmol of ClO4- was produced over

11 days, giving a daily rate of perchlorate ion formation of 4.89 μmol / day. The rates of

perchlorate ion formation were calculated for all samples in each chlorate ion experiment.

To determine the reaction order with respect to the concentration of chlorate ion, the

natural log of the rate of perchlorate ion formation (measured experimentally) was plotted

versus the natural log of the chlorate ion concentration in solutions with the same

concentration of hypochlorite ion. Figure 36 shows the fitted least-squares lines of natural

log of the rate of perchlorate ion formation versus natural log of chlorate ion

concentration over different incubation periods at 30 ºC.

-12.5

-12.0

-11.5

-11.0

-10.5

-10.0

-9.5

-9.0

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0ln[ClO3

- ]

y = 1.067x - 10.37R2 = 0.999

y = 0.984x - 10.44R2 = 1.000

y = 1.030x - 10.35R2 = 0.999

Days 0-11

Days 11-21

Days 21-34ln( d

[ClO

4- ]/dt

)

Figure 36. Fitted natural log lines of the rate of perchlorate ion formation as a

function of chlorate ion concentration at 30 ºC in solutions at constant hypochlorite ion at pH ~12.5

Because the rate of hypochlorite ion decomposition increases with temperature,

differences in decomposition rates of solutions at constant hypochlorite ion and at various

concentration of chlorate ion were observed over shorter incubation periods. Figure 34

(b) and Figure 35 (b) shows that hypochlorite ion solutions with higher concentration of

chlorate ion decompose faster. This was taken into account, and the rate of perchlorate

ion formation as a function of chlorate ion concentration was fitted for hypochlorite ion

solutions with a difference in concentration of hypochlorite ion of less than 5 %. Figures

37 and 38 show fitted plots of the ln (Rate) versus ln [ClO3-] for chlorate ion spiked

experiments at 40° and 50 ºC, respectively.

68

y = 1.007x - 9.24R2 = 0.998

y = 1.083x - 9.41R2= 0.985

-11.0

-9.0

-7.0

-5.0

-3.0

-1.2 -0.6 0.0 0.6 1.2

Day 0-7

Day 7-14

ln[ClO3- ]

ln( d

[ClO

4- ]/dt

)

Figure 37. Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 40 ºC in solutions at constant hypochlorite ion at pH ~12.5

-3.2

-2.4

-1.6

-0.8

0.0

0.8

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0

ln[ClO3- ]

ln( d

[ClO

4- ]/dt

)

y = 0.981x - 0.354R2 = 0.996

Day 0 to 1

y = 1.044x - 0.568R2 = 0.999

Day 1 to 2


function of chlorate ion concentration at 50 ºC in solutions at constant hypochlorite ion at pH ~12.5

The slope of the lines in Figures 36-38 represents the reaction order with respect

to chlorate ion concentration (p), while the intercept is the sum of ln(k2) and m× ln[OCl-]

(Equation 28).

]ln[]ln[ln)ln( 32−− ×+×+= ClOpOClmkRate (28)

69

The slope and R-squared values (R2) for the chlorate ion order-fitting at various

temperatures are summarized in Table 16. As can be observed from Table 16, fitting

ln[Rate] versus ln[ClO3-] produces a linear correlation with an average R2=0.994, and

average slope of 1.035 ± 0.042 for experiments conducted at four temperatures. This

strongly suggests that the reaction order is first order in chlorate ion concentration.

Table 16. Reaction order with respect to chlorate ion and corresponding correlation coefficients in solutions at constant hypochlorite ion at pH ~ 12.5 and various temperatures

T, ºC Slope R2 30 1.067 0.999 30 0.984 1.000 30 1.030 0.998 40 1.007 0.998 40 1.083 0.985 50 0.981 0.996 50 1.044 0.999 75 1.086 0.980

Mean 1.035 0.994 Std. Dev. 0.042 0.008

RSD 4.06 0.81

4.2 Reaction Order with Respect to Chlorate Ion: ln (d[ClO4-]/dt) vs. ln [OCl-]

In this set of experiments, aliquots of concentrated stock hypochlorite ion

solutions, with at least 70 g/L OCl- were diluted, to generate duplicate solutions at

various hypochlorite ion concentrations. Sodium chlorate was added to solutions to keep

chlorate ion constant. Samples were incubated at 30°, 40°, 50°, and 75 ºC.

Figure 39 (a) shows overlaid plots of perchlorate ion formation as a function of

initial chlorate ion concentration; (b) the effects of initial chlorate ion concentrations on

the decomposition of hypochlorite ion and formation of chlorate ion at 30 ºC. Figures 40

and 41 show the effects of chlorate ion concentration during incubation experiments at 40

ºC and at 50 ºC. Note: the downward-sloping, smoothed lines in Figures 39-41 (b)

represent hypochlorite ion decomposition, and the upward-sloping, smoothed lines show

formation of chlorate ion. For clarity some plots were omitted in Figure 40 (b). As was

70

predicted, the rate of perchlorate ion formation increases with an increase in

concentration of hypochlorite ion.

0

50

100

150

200

250

300

350

0 20 40 60 80 100 120 140 160 180 200Days

ClO

4-m

g/L

Form

ed

83g/L OCl- + 63g/L ClO3-



0

20

40

60

80

100

0 25 50 75 100 125 150 175 200O

Cl-

g/L

Dec

ompo

sed

0

20

40

60

80

100

ClO

3-g/

L Fo

rmed





hypochlorite ion and formation of chlorate ion at 30 ºC in solutions at various initial concentrations of hypochlorite ion at pH ~12.5

050

100150200250300350

0 7 14 21 28 35 42 49 56 63 70 77Days

ClO

4-m

g/L

Form

ed








0102030405060708090

100

0 5 10 15 20 25 30 35Days

g/L

OC

l-D

ecom

pose

d

0102030405060708090100

g/L

ClO

3-Fo

rmed

70 g/L OCl- + 51 g/L ClO3-

52 g/L OCl- + 52 g/L ClO3-

45 g/L OCl- + 51 g/L ClO3-

25 g/L OCl- + 52 g/L ClO3-



(b) (a)

(b) (a)

71

0

50

100

150

200

250

300

350

0 1 2 3 4 5 6 7 8 9 10Days

ClO

4-m

g/L

Form

ed







g/L

OC

l-D

ecom

pose

d

g/L

ClO

3-Fo

rmed

0

20

40

60

80

100

0 1 2 3 4 5 6 7 8 9 10Days

0

20

40

60

80

100



To determine the reaction order with respect to hypochlorite ion, the

concentration of chlorate ion was constant. Solutions at various initial concentrations of

hypochlorite ion decomposed differently over longer periods, as was shown during the

chlorate ion spike experiment. Table 17 shows changes in the concentration of chlorate

ion in solutions with various concentration of hypochlorite ion incubated at 30 ºC. Thus,

to minimize the error in the measurement of the rate of perchlorate ion formation due to

increasing variation in chlorate ion over time, the order with respect to hypochlorite ion

was determined based on the data collected during the first 34 days. Table 17. Decomposition of hypochlorite ion at 30 ºC in solutions at constant

chlorate ion concentration at pH ~12.5

Days [OCl-]0 1.594 M

[OCl-]0 0.982 M

[OCl-]0 0.206 M Mean

Std. Dev. RSD

0 0.756 0.780 0.774 0.770 0.012 1.62 11 0.851 0.802 0.806 0.820 0.027 3.32 21 0.899 0.820 0.794 0.838 0.055 6.53 34 0.923 0.834 0.796 0.851 0.065 7.66 49 1.026 0.891 0.826 0.914 0.102 11.16 63 1.077 0.898 0.822 0.932 0.131 14.04

(b) (a)

72

Figures 42-45 show linear plots of the natural log of the rate of perchlorate ion

formation versus the natural log of hypochlorite ion, at different temperatures. The

observation from these figures is that, in general, the slope of the least squares line is

consistently above 1.0.

-15.0

-14.0

-13.0

-12.0

-11.0

-10.0

-9.0

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0

ln[OCl- ]

ln( d

[ClO

4- ]/dt

)y = 1.04x - 11.3R2 = 0.961

y = 1.62x - 11.2R2 = 0.998

y = 1.87x - 11.0R2 = 0.996

Days 0-11

Days 11-21

Days 21-34


function of hypochlorite ion concentration at 30 ºC in solutions at constant chlorate ion at pH ~12.5

-14-12-10-8-6-4-20

-2.0 -1.5 -1.0 -0.5 0.0 0.5

y = 1.14x - 7.95R2 = 0.979

y = 1.72x - 9.99R2 = 0.985

y = 1.67x - 9.87R2 = 0.923

Day 0-7

Day 7-14

Day 14-21

ln[OCl- ]

ln( d

[ClO

4- ]/dt

)


function of hypochlorite ion concentration at 40 ºC in solutions constant in chlorate ion, with pH ~12.5

73

-6.0

-5.0

-4.0

-3.0

-2.0

-1.0

0.0

-2.0 -1.5 -1.0 -0.5 0.0 0.5 1.0

Day 0 to 1

Day 1 to 2

Day 2 to 3

y = 1.45x - 1.54R2 = 0.991

y = 1.25x - 2.01R2 = 0.864

y = 1.18x - 1.32R2 = 0. 986

ln[OCl- ]

ln( d

[ClO

4- ]/dt

)



-9.0-8.0-7.0-6.0-5.0-4.0-3.0-2.0-1.00.0

-2.0 -1.5 -1.0 -0.5 0.0 0.5

Day 0 to 1 y = 1.27x - 6.13R2 = 0.969

ln[OCl- ]

ln( d

[ClO

4- ]/dt

)



The fact that the reaction order with respect to hypochlorite ion appears to vary

markedly and consistently above one suggests that there is another unconsidered variable

involved. To diagnose whether the order with respect to hypochlorite ion is higher than

one, a separate set of experiments at various temperatures with a constant molar product

([OCl-]×[ClO3-] = molar product) were conducted. In this set of experiments, solutions

74

with various concentrations of hypochlorite ion and chlorate ion but at constant molar

product theoretically would yield the same rate of perchlorate ion formation, if the

reaction is first order with respect to either chlorate ion or hypochlorite ion. Figures 46

and 47 show perchlorate ion formation, decomposition of hypochlorite ion, and formation

of chlorate ion as a function of time at 30° and 50 ºC. Note: the downward sloping

smoothed lines and the upward-sloping curves in Figures 46-47 (b) represent

hypochlorite ion decomposition and chlorate ion formation, respectively.

0

20

40

60

80

100

120

140

0 20 40 60 80 100 120 140 160 180 200Days

ClO

4-m

g/L

Form

ed

1.64 M OCl- + 0.18 M ClO3-

0.99 M OCl- + 0.30 M ClO3-

0.20 M OCl- + 1.49 M ClO3-

00.20.40.60.81.01.21.41.61.8

0 25 50 75 100 125 150 175 200Days

00.20.40.60.81.01.21.41.6

[OC

l- ]

1.64 M OCl- + 0.18 M ClO3-

0.99 M OCl- + 0.30 M ClO3-

0.20 M OCl- + 1.49 M ClO3-

[ClO

3- ]


hypochlorite ion and formation of chlorate ion at 30 ºC in solutions at constant molar product at pH ~12.5

As is evident from Figures 46 and 47, solutions containing the highest

concentration of hypochlorite ion had significantly higher rates of the perchlorate ion

formation. Thus, it is proposed that either the rate of perchlorate ion formation is not

second-order, but rather a higher order process due to larger dependence on concentration

of hypochlorite ion, or there is an experimental variable not accounted for.

(b) (a)

75

0

20

40

60

80

100

120

140

0 1 2 3 4 5 6 7 8 9 10Days

ClO

4-m

g/L

Form

ed

1.68 M OCl- + 0.16 M ClO3-

1.01 M OCl- + 0.26 M ClO3-

0.20 M OCl- + 1.32 M ClO3-

00.20.40.60.8

11.21.41.61.8

0 1 2 3 4 5 6 7 8 9 10Days

[OC

l- ]

0

0.2

0.4

0.6

0.8

1

1.2

1.4

[ClO

3- ]

1.68 M OCl- + 0.16 M ClO3-

1.01 M OCl- + 0.26 M ClO3-

0.20 M OCl- + 1.32 M ClO3-


hypochlorite ion and formation of chlorate ion at 50 ºC in solutions at constant molar product at pH ~12.5

4.3 Multiple Reaction Pathways

To investigate the discrepancy in the order with respect to hypochlorite ion,

multiple reaction pathways were considered. Fitting of the data, based on the assumption

that perchlorate ion formation is a second-order overall process, resulted in variable

reaction order with respect to hypochlorite ion. This suggested an order higher than one.

Thus, a second reaction involving hypochlorite ion was considered. Thus, hypothetically,

if perchlorate ion formation is a two-reaction process, then the reactions are either

parallel or consecutive pathways.

A parallel reaction pathway, given by Equation 28, may involve a sum of

reactions that are first-order and second-order in hypochlorite ion. Where k1 is a second-

order rate constant, and k2 is a third-order rate constant.

][][]][[][3

2231

4 −−−−−

+== ClOOClkClOOClkdt

ClOdRate (28)

A consecutive reaction pathway, given by Equation 29, may involve a

competitive pre-equilibrium reaction with another hypochlorite ion to form an

(b) (a)

76

intermediate. This pre-equilibrium reaction is followed by decomposition of the

intermediate to form either reactants or products. Where ka, is a rate constant that drives

the reaction to form products, and the value of rate constant, kb, determines the extent of

the forward reaction.

][1

][][][ 32

4−

−−−

+==

OClkClOOClk

dtClOdRate

b

a (29)

4.3.1 Parallel Reaction Pathway

To investigate the parallel reaction pathway, data from experiments that varied in

hypochlorite ion but constant in chlorate ion concentration were used to determine the

values of k1 and k2. If Equation 28 is rearranged by dividing by the concentration of

hypochlorite ion and concentration of chlorate ion (molar product of the two reactants),

Equation 30 is obtained.

][]][[

213

−−−

+= OClkkClOOCl

Rate (30)

When the [Rate / ([OCl-]× [ClO3-])] is plotted versus [OCl-] (for constant chlorate

ion experiments), a linear correlation will provide values of k2 (the slope of the line) and

k1 (the intercept). Figure 48 shows linear plots at different temperatures. Table 18 shows

the values of k2 and k1 at each temperature, calculated from the slope and intercept of the

fitted lines in Figure 48.

77

y = 8.42∗10-6x + 7.29∗10-6

R2 = 0.995

0.0

5.0 ∗10-5

1.0 ∗10-5

1.5 ∗10-5

2.0 ∗10-5

2.5 ∗10-5

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8[OCl-]

Rat

e/[O

Cl- ][

ClO

3- ]

y = 3.47 ∗10-5x + 1.87∗10-5

R2 = 0.984

0.01.0∗10-52.0∗10-53.0∗10-54.0∗10-55.0∗10-56.0∗10-57.0∗10-5

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

Rat

e/[O

Cl- ][

ClO

3- ]

[OCl-]

y = 1.32∗10-4x + 8.63∗10-5

R2 = 0.999

0.05.0∗10-41.0∗10-41.5∗10-42.0∗10-42.5∗10-43.0∗10-43.5∗10-4

0.0 0.5 1.0 1.5 2.0

[OCl-]

Rat

e/[O

Cl- ][

ClO

3- ]

Figure 48. Parallel reaction pathway linear fitted plots of (a) 30 ºC experiment;

(b) 40 ºC experiment; (c) 50 ºC experiment; in solutions at constant chlorate ion and various inial voncentrations of hypochlorite ion at pH ~12.5

(a) 30 ºC

(b) 40 ºC

(c) 50 ºC

78

Table 18. Parallel reaction pathway experimental rate constants in solutions, at various hypochlorite ion and constant chlorate ion, at pH ~12.5

T, ºC k2 × 106 M-2·d-1 k1 ×106

M-1·d-1 R2 k2/k1 30 8.43 7.29 0.995 1.16 40 34.7 18.7 0.984 1.85 50 132 86.3 0.999 1.53

An initial analysis of k2 and k1 values shows that both rate constants increase with

temperature; however, the ratio of k2/k1 is changing. In general, it may be expected that if

there is a temperature dependence, the effect of temperature on the two rate constants

may not be the same. The ratio of the rate constants therefore would be expected to

increase or decrease (have a consistent trend) but not go up and down. This suggests that

although a reasonable fitting can be established, there still appears to be a degree of

uncertainty in the parallel pathway model.

4.3.2 Consecutive Reaction Pathway

To investigate the consecutive reaction pathway, data from experiments that

varied in hypochlorite ion concentration at constant chlorate ion concentration were used

to determine the values of ka and kb. Equation 29 can be rearranged by dividing both

sides of the equation by the concentration of hypochlorite ion to obtain Equation 31:

]])[[1(

][][][

32

−−

−−

− +=

OClOClkClOOClk

OClRate

b

a (31)

If, Equation 31 is inverted and both sides are multiplied by the concentration of

chlorate ion, Equation 32 results:

][][

]][])[[1(]][[

32

33−−

−−−−− +=

ClOOClkClOOClOClk

RateClOOCl

a

b (32)

Equation 33 results from cancelling the common terms on the right side of

Equation 32.

a

b

a kk

OClkRateClOOCl

+=−

−−

][1]][[ 3 (33)

By plotting the quantity [OCl-]× [ClO3-]/(Rate) versus 1/[OCl-], the values of ka

and kb can be determined from the least-squares lines: the slope equals 1/ka and the

79

intercept is kb/ka (From Equation 33). Figure 49 shows the fitting of the data from

constant chlorate ion and variable hypochlorite ion experiments at different temperatures.

Table 19 shows the values of ka and kb at each temperature, calculated from the slope and

intercept of the fitted lines in Figure 49.

y = 1.3∗104x + 4.6∗104

R2 = 0.951

0.0

0.2 ∗105

0.4 ∗105

0.6 ∗105

0.8 ∗105

1.0 ∗105

1.2 ∗105

0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5 5.0

1/[OCl-]

[OC

l- ][C

lO3- ]

/ Rat

e

0.05.0∗104

1.0∗104

1.5∗104

2.0∗104

2.5∗104

3.0∗104

3.5∗104

0 0.3 0.6 0.9 1.2 1.5 1.8 2.11/[OCl-]

y = 1.0∗104x + 8.32∗103

R2 = 0.985

[OC

l- ][C

lO3- ]

/ Rat

e

y = 1.7∗103x + 2.7∗103

R2 = 0.997

0.0

0.2∗104

0.4∗104

0.6∗104

0.8∗104

1.0∗104

1.2∗104

0.0 1.0 2.0 3.0 4.0 5.0

1/[OCl-]

[OC

l- ][C

lO3- ]

/ Rat

e

Figure 49. Consecutive reaction pathway linear fitted plots of (a) 30 ºC

experiment; (b) 40 ºC experiment; (c) 50 ºC experiment; in solutions at constant chlorate ion and various hypochlorite ion at pH ~12.5

(a) 30 ºC

(b) 40 ºC

(c) 50 ºC

80

Table 19. Consecutive reaction pathway experimental rate constants at various initial concentrations of hypochlorite ion and constant chlorate ion at pH ~12.5 T, ºC ka ×106

M-2·d-1 kb M-1·d-1 R2 kb / ka 30 76.1 3.49 0.951 45,800 40 100 0.83 0.985 8,300 50 591 1.59 0.997 2,700

The linear fit of the experimental data to a consecutive reaction pathway for the

30 ºC experiment, shown in Figure 49, has a low R-squared value of 0.951. In addition,

as shown in Table 19, the second-order rate constant, kb, is highly-variable at different

temperatures, with no consistent trend. This would indicate that either more experiments

are needed to provide better fits or the model based on the consecutive reaction pathway

does not fit the experimental data.

The data analysis in this and previous sections, demonstrates the possibility that

the reaction order with respect to hypochlorite ion may be greater than one; however,

both models based on a parallel or consecutive reaction pathway fit the experimental data

poorly and provide inconsistent trends in the values of determined rate constants. This

strongly indicates a dependence of the rate constant(s) on an additional variable that has

not been considered in attempted models to fit the experimental data.

4.4 Ionic Strength Effect on the Rate of Perchlorate Ion Formation

Based on the analysis of data and observations drawn from experiments, it was

hypothesized that perhaps a simpler explanation for the variability based on the reaction

order for hypochlorite ion may arise from the rate dependence on the ionic strength of

perchlorate ion formation.

In Chapter 3, an increase in ionic strength by addition of sodium chloride to the

hypochlorite solutions resulted in an increased rate of perchlorate ion formation (Figure

28 page 56). When examining the measured ionic strength of solutions with various

concentrations of chlorate ion as presented in Table 20, it was discovered that the

differences in ionic strength of these solutions were not significant. Thus, the

experimentally-measured reaction order with respect to chlorate ion of one must be valid.

81

Table 20. Ionic strength (μ) of hypochlorite ion solutions at various chlorate ion at 40 ºC experiments (TDS = Total Dissolved Solids)

[ClO3-], mol/L [OCl-], mol/L μ, mol/L pH TDS, g/L

1.788 1.355 7.96 12.99 318 1.449 1.384 7.66 12.52 307 1.254 1.395 7.48 12.96 299 0.907 1.393 7.11 12.95 284 0.432 1.372 6.56 12.93 263 0.337 1.370 6.46 12.94 259 Mean 1.38 7.21 12.88

Std. Dev. 0.02 0.60 0.18 RSD 1.11 8.37 1.38

Incubation experiments at 40 ºC had the largest number of sample solutions with

varying hypochlorite ion. The ionic strength and pH values are presented in Table 21.

What is evident from Table 21 is the change in ionic strength of the solutions with

various concentrations of hypochlorite ion is significant (1.76-6.72 mol/L).

Table 21. Ionic strength (μ) of hypochlorite ion solutions at various hypochlorite ion at 40 ºC experiments (TDS=Total Dissolved Solids)

[OCl-], mol/L [ClO3-], mol/L μ, mol/L pH TDS, g/L

1.366 0.619 6.72 13.00 268.7 1.371 0.610 6.80 12.99 272.0 1.166 0.619 6.02 12.96 240.6 1.012 0.623 5.38 12.91 215.1 0.871 0.605 4.80 12.84 192.1 0.867 0.615 4.78 12.82 191.3 0.646 0.618 3.78 12.70 151.4 0.478 0.621 3.14 12.60 125.7 0.485 0.614 3.10 12.56 124.1 0.178 0.610 1.76 12.21 70.2 Mean 0.620 4.63 12.76

Std. Dev. 0.006 1.67 0.25 RSD 0.90 36.09 1.95

In fact, the observed rates of perchlorate ion formation seem to increase with the

increase in ionic strength and coincidently with the increase in hypochlorite ion

concentration, shown in Figure 50.

82

y = 1.24∗10-6x2 – 6.44 ∗10-8x -1.37∗10-6

R2 = 0.999

0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.00.0

2.0∗10-5

4.0∗10-5

6.0∗10-5

ln( d

[ClO

4- ]/dt

)

Ionic Strength (μ), mol/L

0.0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

y = 1.78∗10-5x2 + 1.70 ∗10-5x -1.63∗10-6

R2 = 0.997

0.0

2.0∗10-5

4.0∗10-5

6.0∗10-5

ln( d

[ClO

4- ]/dt

)

[OCl-], mol/L Figure 50. Rate of perchlorate ion formation as a function of (a) ionic strength;

(b) concentration of hypochlorite ion, at 40 ºC

The plots of the rate of perchlorate formation as a function of ionic strength and

hypochlorite ion show very similar behavior, as can be seen in Figure 50. At low ionic

strength and low concentrations of hypochlorite ion, the rate of perchlorate ion formation

changes more slowly. At ionic strength above 2.0 M and 0.5 M OCl-, the rate of

perchlorate ion formation correlates linearly. The marked agreement between both plots

is a strong indication that both ionic strength and hypochlorite ion concentration affects

the rate of perchlorate ion formation and needs to be deconvoluted.

(a)

(b)

83

The main objectives are to elucidate the rate law for formation of perchlorate ion

and determine the rate constant(s) and provide a readily usable model (with the fewest

number of parameters) for perchlorate ion formation. A strong dependence of the rate of

perchlorate ion formation on the ionic strength, introduces an additional term that must be

related to the rate constant(s). To investigate the relationship of the ionic strength and the

rate constant(s) of perchlorate ion formation, an assumption about the rate law is needed.

Multiple reactions have been considered but involve two rate constants. A

simpler starting point can be based on a single, second-order reaction that takes place

between hypochlorite and chlorate ions. The rate of perchlorate ion formation is shown in

Equation 27:

pm ClOOClkdt

ClOdRate ][][][32

4 −−−

×== (27)

The order with respect to chlorate ion has been determined to be first-order, based

on experimental data from solutions at various chlorate ion and constant ionic strength

(Tables 16 and 21). Because the solutions at various concentration of the hypochlorite

ion were coincidently at various ionic strengths, a hypothesis is needed for differentiation

of the perchlorate ion formation dependence on hypochlorite ion and ionic strength. As

an approach to deconvolute both effects, the order with respect to hypochlorite ion was

assumed to be first-order. Thus, a simpler model based on reaction that is first-order in

both (m, p = 1) chlorate and hypochlorite ions is obtained, allowing determination of the

second-order rate constant by Equation 34.

][ClO][OCl

Ratek3

2 −− ×= (34)

The fitting of the experimental rate of perchlorate ion formation as a function of

ionic strength from the 40 ºC experiments, as shown in Figure 50, demonstrated linearity

at ionic strengths greater than 2 M. However, before any further conclusions can be

made, an expression relating the second-order rate constant, k2, and ionic strength is

needed.

84

4.4.1 Dependence of the Second-Order Rate Constant on the Ionic Strength

Transition state theory (developed by Henry Eyring91) describes reaction rates

based on formation of an activated complex. For a biomolecular reaction between

hypochlorite ion and chlorate ion, the formation of activated complex and products can

be described by Equation 35.

−−−−−− +⎯→⎯⎯→⎯+ ClClO]ClOOCl[ClOOCl 4‡

33‡

2 kk (35)

Where, k2, is the experimental second-order rate constant, and [OCl-ClO3-]‡

denotes the activated complex. The formation of the quasi equilibrium between reactants

and [OCl-ClO3-]‡ intermediate can be described by an equilibrium constant K‡, shown in

Equation 36:

−−−−

−−

×=

33

‡‡3‡

][ClO][OCl

]ClOOCl[ClOOCl

Kγγ

γ (36)

where, γ is the activity coefficient. Thus the rate of this reaction can be defined as shown

in Equation 37.

‡3

‡32

4 ]ClOOCl[][][][ −−−−−

=×== kClOOClkdt

ClOdRate (37)

By defining the quantity [OCl-ClO3-]‡ using Equation 36 and substituting into

Equation 37, the second-order rate constant, k2, can be expressed by Equation 38.

‡

‡‡2

3

γ

γγ−−

×=ClOOCl

Kkk (38)

where the quantity k‡×K‡, is defined as the rate constant at infinite dilution, and is

commonly defined as kref or ko. Thus Equation 38 really is the same as the Brønsted-

Bjerrum Equation92, shown by Equation 39:

‡2

3

γ

γγ−−

=ClOOCl

okk (39)

Thus, by taking the log of both sides of the Equation 39, Equation 40 results.

91 Espenson, J. H. Chemical Kinetics and Reaction Mechanisms, 2nd ed.; McGraw-Hill, Inc., 1995. 92 Espenson Chemical Kinetics and Reaction Mechanisms.

85

)log()log()log( ‡23

γ

γγ−−

+=ClOOCl

okk (40)

A reduced form of a Debye-Hückel Equation by Güntelbuerg93, shown in

Equation 41, can be used for all electrolytes at 25 ºC to relate the activity coefficient to

the ionic strength (γi denotes activity coefficient for each species i).

μμ

γ+×

−=1

)log(2i

iZA (41)

where Z1 and Z2 are the charges of the electrolyte ions and A is a constant that increases

with temperature.93 The ionic strength of the hypochlorite ion solutions has been

measured in range of 0.8-8M, thus the modified Equation 41 by Gugenheim, which adds

a linear term94 for solutions with ionic strength above 0.1 M, by Equation 42, is needed.

μμμ

γ bZA i

i ++×

−=1

)log(2

(42)

where b is an adjustable parameter95, Davies in his modification of Equation 41 used the

value of 0.2 for b and 0.50 for the term A at 25 ºC.96 By constructing three equations like

Equation 42 and combining them into Equation 40, Equation 43 results:

)1

2()log()log( 3

2 μμ

μb

ZAZkk

ClOOClo −

++=

−−

(43)

The objective is to quantify the effect of the ionic strength on the rate constant by

reducing the correlation to one single equation that can be used to predict changes in rate

constant as a function of the ionic strength. The different approximations of the Debye-

Hückel Equation and derivations of the limiting law by Güntelbuerg, Gugenheim, and

Davies, were based on the experimental data. Thus, the effects of ionic strength on the

rate constant described in Equation 43 need to be examined further.

93 Robinson, R. A. and Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths Publications Limited: London, 1959. 94 Robinson and Stokes Electrolyte Solutions. 95 Harned, H. S. and Owen, B. B. The Physical Chemistry of Electrolytic Solutions, 2nd ed.; Reinhold Publishing Corporation: New York, 1950. 96 Davies, C. W. "397. The extent of dissociation of salts in water. Part VIII. An equation for the mean ionic activity coefficient of an electrolyte in water, and a revision of the dissociation constants of some sulphates" J. Chem. Soc. 1938, 1938, 2093-2098.

86

In the context of chemical reactions, higher ionic strength results in an increase of

the net charge of the ionic atmosphere of each ion, thereby reducing attraction between a

cation and an anion and reducing the repulsive forces between ions of the same polarity.

Thus, ions of the same charge polarity are more likely to come together, as they are

stabilized by electrostatic attractions. An increase in ionic strength favors interactions of

the charged ions and activated complex over those between the reactants and the charged

ions, thus the rate constant increases. Furthermore, the formation of activated complex

due to ion-ion pairing becomes increasingly dependent on ionic strength above 0.01

mol/L and because deviations occur, an additional linear term is incorporated in the

Equation 43 for corrections.97 Previously it was found that the decomposition of

hypochlorite ion was found to be strongly dependent on the ionic strength, due to the ion-

ion interactions.98 Equation 43 incorporates both the effect based on increased activity of

ionic species, described by the term )1/( μμ + and the effect of ion-ion interactions,

described by the term μb .

Both terms were plotted as a function of ionic strength, shown in Figure 51, so

that their relative contributions can be evaluated.

0

0.5

1.0

1.5

2.0

2.5

3.0

0 1 2 3 4 5 6 7 8

μ, mol/L

)1/( μμ +

μb

Rel

ativ

e C

ontr

ibut

ion

Figure 51. Plot of the )1/( μμ + and μb terms as a function of ionic strength.

Note: value of 0.5 was assumed for the b term 97 Robinson and Stokes Electrolyte Solutions. 98 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

87

As can be observed from Figure 51, the contribution of the ion-ion interaction

term99 becomes more significant at μ ≥ 1 M, by assuming b = 0.5. Thus, an empirical

relationship between ionic strength, the second-order rate constant, k2, and the rate

constant at infinite dilution, k0, is approximated by Equation 44:

)log(k)log(k o2 += μb (44)

The values of k2 were calculated for experiments conducted at different

temperatures by using Equation 44 and plotted against the ionic strength of solutions,

shown in Figure 52.

30 ºC: y = 0.0738x - 10.1R2 = 1.000

-10.5

-10.0

-9.5

-9.0

-8.5

-8.0

0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0

μ, mol/L

log(

k 2)

50 ºC: y = 0.0788x - 9.00R2 = 0.978

40 ºC: y = 0.0838x - 9.66R2 = 0.971

Figure 52. Overlaid linear plots of log of second-order rate constant versus ionic

strength in solutions with various initial concentrations of hypochlorite ion at different temperatures

The slopes of the least-squares lines are equal to the b term, and the intercepts are

log of the second-order rate constant at infinite dilution. The log (ko) values are

summarized in Table 22. Reasonable agreement (RSD = 6.4 %) is observed in the slopes

of the lines at various temperatures. This demonstrates that the ionic strength correlates

with the second-order rate constant of perchlorate ion formation. Thus, Equation 46 99 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

88

provides the simplified approximation of the changes in the second-order rate constant as

a function of the ionic strength at different temperatures.

Table 22. Slopes and intercepts of least-squares lines shown in Figure 52 Temperature

(ºC) Slope Intercept (log k0)

ko in M-1s-1 (x 10-12)

30 0.0738 -10.1 77.0 40 0.0838 -9.66 217 50 0.0788 -9.00 990

Mean 0.0788 Std. Dev. 0.0050

RSD 6.35

4.4.2 Dependence of the Second-Order Rate Constant on the Temperature

The rate of perchlorate ion formation is enhanced at elevated temperatures in

solutions with similar initial concentrations of hypochlorite and chlorate ions, as can be

seen from Figure 53. This effect must be related to the temperature dependence of the

second-order rate constant at infinite dilution.

0100200300400500600700

0 10 20 30 40 50 60 70 80 90Days

ClO

4-m

g/L

Form

ed

30 ºC [OCl-]0= 84.2 g/L, [ClO3-]0= 14.8 g/L

40 ºC [OCl-]0= 70.5 g/L, [ClO3-]0= 28.1 g/L

50 ºC [OCl-]0= 85.3 g/L, [ClO3-]0= 13.0 g/L

75 ºC [OCl-]0= 75.0 g/L, [ClO3-]0= 24.2 g/L

Figure 53. Smooth-line plots of the perchlorate ion formation as a function of

time in solutions at similar initial hypochlorite and chlorate ions and various temperatures

89

The Arrhenius and the Eyring equations are used to describe the temperature

dependence of a reaction rate. However, the Eyring equation, which is based on

transition state theory, is used for studying kinetics of reactions occurring in liquids. The

Eyring equation100 is shown in Equation 47 below:

)()(

‡‡

RTH

RS

bo eeT

hk

kΔ

−Δ

×××= (45)

ko = Calculated using Equation 27 M-1 · s-1 kb = Boltzmann constant 1.381x10-23 J · K-1 h = Planck constant 6.626x10-34 J · s R = Real Gas constant 8.3145 J/mol · K ΔS‡ = Entropy of activation J/mol · K ΔH‡ = Enthalpy of activation kJ/mol T = Temperature Kelvin (K) By dividing both sides of Equation 45 by T and taking the natural log function of

both sides, Equation 46 results:

RTH

RS

hk

Tk bo

‡‡)ln()ln( Δ

−Δ

+= (46)

To determine ΔH‡ and ΔS‡, ln(ko/T) is plotted as a function of 1/T. The slope of

the linear fit is (-ΔH‡ / R), while the intercept equals the sum of constants and entropy

term [ln (kb / h) + (ΔS‡ / R)]. Temperatures in Table 20 were converted from ºC to

Kelvin to determine values of ln(ko/T) and 1/T terms. Figure 54 shows a plot of

experimentally determined ln(ko/T) vs 1/T.

100 Espenson Chemical Kinetics and Reaction Mechanisms.

90

y = -12165x + 11.038R2 = 0.983

-30

-29

-28

-27

-26

-25

0.00300 0.00315 0.00330 0.003451/T, K

ln(k

o/T)

Figure 54. Linear plot of ln(ko/T) as a function of (1/T)

Based on three points the R2 of the line is 0.983. The relatively low r-squared

correlation coefficient of 0.983 is an indication that more temperatures would result in

more information. However, based on the low variation observed in the fitted slopes of

6.4 %, shown in Table 20, the R2 correlation coefficient of 0.983 is satisfactory and

provides a simple model that does not introduce additional power terms.

Thus, the approximate values of thermodynamic activation parameters for the

formation of perchlorate ion in bulk, concentrated hypochlorite ion solutions at infinite

dilution are ΔH‡ = 101 kJ/mol and ΔS‡= -106 J/mol·K. A large value of the ΔH‡ is

indicative of a slow reaction and a large, negative value of ΔS‡ reflects loss of entropy

from the union of hypochlorite and chlorate ions into a single molecule.101

By substituting the calculated values of ΔH‡ and ΔS‡, based on slope and

intercept values shown in Figure 54, into Equation 45, a generalized expression is

obtained relating second-order rate constant at infinite dilution and thermodynamic

activation parameters, shown in Equation 47.

R106

RT1.01x10

10 eeT102.084

5 −−

××××=ok (47)

101 Espenson Chemical Kinetics and Reaction Mechanisms.

91

4.4.3 Combining the Effects of the Ionic Strength and Temperature on the

Second-Order Rate Constant

The results of the previous sections in this chapter can be summarized with

several key conclusions. The rate of perchlorate ion formation has been shown to

correlate with first-order concentrations of hypochlorite and chlorate ions, ionic strength,

and temperature. The second-order rate constant was determined to increase as a

function of the ionic strength. Temperature dependence of the rate constant at infinite

dilution has also been determined. Thus, by combining the experimentally determined b

value of 0.0788 and Equation 47 into Equation 44, the combined effects of ionic strength

and temperature are given by Equation 48:

)eeT10log(2.084)0.0788()log(k R106

RT1.01x10

102

5 −−

××××+= μ (48)

The second-order rate constant of perchlorate ion formation for solutions of any

ionic strength in the range of 1.9-6.9 mol/L and temperature range 30-50 ºC can be

calculated by using Equation 48. Table 23 shows the experimentally observed values of

second-order rate constant, k2, and predicted values by using Equation 48.

Table 23. Experimental and predicted second-order rate constants at variable ionic strength and temperature (kexp = experimental k2; kpred = predicted k2)

T (ºC)

μ (M)

kexp (M-1 d-1 ×106)

kpred (M-1 d-1 ×106)

% Error

Average % Error

R2

log(k2) vs. μ 30 6.87 21.38 21.76 1.8 30 4.73 14.88 14.76 0.8 30 1.93 9.24 8.88 3.8 2.1 0.9999 40 6.72 63.70 78.65 23.5 40 6.80 65.51 79.82 21.9 40 6.02 61.29 69.24 13.0 40 5.38 55.85 61.67 10.4 40 4.80 49.99 55.55 11.1 40 4.78 48.59 55.35 13.9 40 3.78 40.26 46.19 14.7 40 3.14 35.52 41.11 15.7 40 3.10 34.19 40.81 19.4 16 0.9709 50 6.89 285.56 277.98 2.7 50 4.77 219.79 189.46 13.8 50 1.92 117.24 112.84 3.8 6.7 0.9784

92

The difference between the experimentally-observed and predicted value of k2 is

reported as percent error in Table 23. Given the relatively low R2 values, the average %

errors are acceptable for the hypochlorite ion solutions used in the data fitting. The

predicted values use a minimum number of parameters and provide a simple approach to

the determination of the rate law of perchlorate ion formation based on the bimolecular

reaction between hypochlorite and chlorate ion.

4.5 Conclusions

Multiple reaction models were used to fit the experimental data. The data

analysis revealed that the rate law for perchlorate ion formation can be described as first

order in hypochlorite ion and chlorate ion and a single second-order rate constant with a

strong dependence on the ionic strength. The dependence of the second-order rate

constant on the ionic strength was quantified using a simplified extension of the Debye-

Hückel Equation. The experimental rate constants at infinite dilutions were related to

temperature and thermodynamic parameters determined using the Eyring equation. Thus,

a quantitative expression between identified effects of ionic strength and temperature on

the second-order rate constant is provided to utilities to predict how to decrease the

amount of perchlorate ion formed during storage of hypochlorite ion solutions.

93

CHAPTER 5. THE PERCHLORATE ION FORMATION MODEL:

VALIDATION AND APPLICATIONS

This chapter describes the validation and application of a perchlorate ion

formation model. There are several objectives for this chapter. The first objective is to

validate the established relationship between the second-order rate constant, ionic

strength, and temperature (Equation 50) with the use of hypochlorite ion decomposition

Bleach 2001 to predict perchlorate ion formation in bulk sodium hypochlorite solutions.

The objective for the predictive model is to have an average error of ± 10 %. The

developed model was used to predict the rate of perchlorate ion formation in survey bulk

hypochlorite ion solutions from several utilities in United States, and compared to the

experimental data.

The second objective is to use the model to assist water utilities to assess the

formation of perchlorate ion at different conditions and to discuss the implications of

storing concentrated (undiluted) hypochlorite ion solutions. In addition, the developed

predictive perchlorate ion formation model is applied to survey OSG sodium

hypochlorite solutions and calcium hypochlorite solutions. Finally, the potential

contributions from different sources of hypochlorite ion to perchlorate ion contamination

of drinking water are discussed.

The hypochlorite ion decomposition model, Bleach 2001, was developed as a

predictive tool to guide utilities to develop optimum storage conditions and be used to

predict when the hypochlorite ion solutions are no longer usable. During the

development and validation of Bleach 2001, multiple commercial hypochlorite ion

solutions were tested, and an average error of ± 5 % was reported between

experimentally measured decomposition and predicted decomposition.102 In the design

of the experiments conducted in the current work, Bleach 2001 was used to predict

decomposition of hypochlorite ion and formation of chlorate ion.

Figure 55 shows the decomposition of hypochlorite ion (downward-sloping

curves) and formation of chlorate ion (upward-sloping curves) measured experimentally

and predicted by Bleach 2001 at 30º, 40º, and 50 ºC. 102 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite.

94

102030405060708090

0 25 50 75 100 125 150 175 200Days

50

60

70

80

90

100

110

OC

l-g/

L D

ecom

pose

d

ClO

3-g/

L Fo

rmed

ExperimentBleach 2001

0102030405060708090

0 5 10 15 20 25 30 35Days

0102030405060708090Bleach 2001

Experiment

OC

l- g/L

Dec

ompo

sed

ClO

3-g/

L Fo

rmed

0

102030405060708090

100

0 2 4 6 8 1050556065707580859095100

Days

ExperimentBleach 2001

OC

l- g/L

Dec

ompo

sed

ClO

3-g/

L Fo

rmed

Figure 55. Smoothed-line plots of hypochlorite ion decomposition and chlorate

ion formation determined experimentally in conjunction with Bleach 2001 (Error bars set at ± 10 %), for solutions at (a) [OCl-]0 = 82 g/L, [ClO3

-]0 = 63 g/L at 30 ºC; (b) [OCl-]0 = 70 g/L, [ClO3-]0 = 51 g/L at 40

ºC; (c) [OCl-]0 = 83 g/L, [ClO3-]0 = 50 g/L at 50 ºC

(a) 30 ºC

(b) 40 ºC

(c) 50 ºC

95

The average % difference between measured hypochlorite ion and predicted

values by Bleach 2001, at 30 ºC over a 200-day period, was 8.8 % ± 4.2 (n = 7); at 40 ºC

over 34-day period, the average % difference was 2.2 % ± 8.8 (n = 4); and at 50 ºC, the

average % difference over 10 days was 6.6 % ± 4.6 (n = 6). Thus, the predictive

expression for the second-order rate constant (Equation 50) was used with Bleach 2001,

to predict the formation of chlorate and perchlorate ions and the decomposition of

hypochlorite ion at various temperatures, ionic strengths, and initial concentrations of

hypochlorite and chlorate ions.

5.1 Predicted Perchlorate Ion Formation in Bulk Sodium Hypochlorite Solutions

The bulk hypochlorite ion solutions used in various incubation experiments and

eventually for the determination of the perchlorate ion formation rate law were chosen for

validation of the predictive expression of the second-order rate constant. Based on the

ionic strength of the hypochlorite ion solutions and the incubation temperature, second-

order rate constants were generated, using Equation 48.

)eeT10log(2.084)0.0788()log(k R106

RT1.01x10

102

5 −−

××××+= μ (48)

The Bleach 2001 model was used to predict the decomposition of hypochlorite

ion and formation of chlorate ion at various temperatures. The predicted values of k2, and

the predicted concentrations of hypochlorite and chlorate ions were used to generate the

rates of perchlorate ion formation for each solution, using Equation 27.

][][][32

4 −−−

×== ClOOClkdt

ClOdRate (27)

As an example, to simulate the perchlorate ion formation at 70.4 g/L OCl-, 50.1

g/L ClO3-, and μ = 6.76 mol/L during a 40 ºC incubation, the value of k2 was calculated

by entering temperature (converted from 40 ºC to Kelvin) and ionic strength into

Equation 50. The obtained value of k2, 9.15×10-10 L·mol-1·s-1, was converted to units of

L·mol-1·d-1, producing a value of 7.91×10-5. The predicted value of k2, and predicted

changes in concentration of hypochlorite and chlorate ions by Bleach 2001 were

96

multiplied to calculated the rate of perchlorate ion formation as a function of

concentration of hypochlorite and chlorate ions. The obtained rates were multiplied by

time, in days, and an incremental increase in perchlorate ion concentration was calculated

per time interval shown in Table 24.

Table 24. Predicted changes in hypochlorite ion, chlorate ion, and d[ClO4-]/dt as

a function of time at 40 ºC

Day [OCl-], (mol/L)

[ClO3-],

(mol/L) d[ClO4

-]/dt, (mol/L/day ×106)

d[ClO4-]/dt,

(mg/L / day) mg/L ClO4

- produced

0 1.369 0.600 65.0 6.5 6.46 1 1.310 0.619 64.1 6.4 6.37 2 1.256 0.635 63.1 6.3 6.27 3 1.206 0.651 62.1 6.2 6.17 4 1.160 0.665 61.0 6.1 6.07 5 1.118 0.678 59.9 6.0 5.96 6 1.078 0.690 58.9 5.9 5.85 7 1.042 0.701 57.8 5.8 5.75 14 0.841 0.764 50.8 5.1 35.3 21 0.705 0.806 44.9 4.5 31.3 28 0.607 0.836 40.1 4.0 27.9 35 0.533 0.859 36.2 3.6 25.2

The summation of incremental increases in perchlorate ion concentration over

initial concentration as a function of time is recorded at each day-interval. Figure 56

shows overlaid plots of predicted perchlorate ion concentration as a function of time and

experimentally measured concentration. The smoothed-lines represent predicted changes

in perchlorate ion formation, with error bars set at a fixed ± 10 %.

97

0

50

100

150

200

250

300

350

0 50 100 150 200

Days

ClO

4- mg/

L

[OCl-]0= 83 g/L; [ClO3-]0=50 g/L

[OCl-]0= 52 g/L; [ClO3-]0=50 g/L

[OCl-]0= 10 g/L; [ClO3-]0=50 g/L

0

50

100

150

200

250

300

350

0 7 14 21 28 35 42 49 56 63 70 77Days

[OCl-]0= 60 g/L; [ClO3-]0=50 g/L

[OCl-]0= 33 g/L; [ClO3-]0=50 g/L

ClO

4- mg/

L

[OCl-]0= 70 g/L; [ClO3-]0=50 g/L

0

50

100

150

200

250

300

350

0 2 4 6 8 10

Days

[OCl-]0= 83 g/L; [ClO3-]0=50 g/L

[OCl-]0= 52 g/L; [ClO3-]0=50 g/L

[OCl-]0= 10 g/L; [ClO3-]0=50 g/L

ClO

4-m

g/L

Figure 56. Overlaid smoothed-line plots of predicted (Error bars set at ± 10 %)

perchlorate ion formation and determined experimentally, for solutions incubated at (a) 30 ºC; (b) 40 ºC; (c) 50 ºC

(a) 30 ºC

(b) 40 ºC

(c) 50 ºC

98

As can be seen from Figure 56, in general the experimental and predicted

formations of perchlorate ion agree for solutions with various concentrations of

hypochlorite ion stored at different temperatures. On average the initial predicted rates

are well within ± 10 %, and the model is clearly capable of predicting the formation of

perchlorate ion as a function of time. Higher deviations in the predicted concentration of

perchlorate ion were observed in solutions stored at 50 ºC after several days. To test the

model further, several hypochlorite ion solutions were acquired from utilities in United

States and incubated at 50 ºC. These results are discussed in the next section.

5.2 Predicted Perchlorate Ion Formation in Real-World Bulk Sodium

Hypochlorite Solutions

Bulk hypochlorite ion solutions were obtained from five individual utilities

located in AZ, CA, FL, GA, and OH. Samples of raw water, finished water, and

distribution system samples were also collected. Hypochlorite ion solutions were

collected in duplicate, including a sample, quenched with malonic acid, and non-

quenched sample that was cooled to 4 ºC. The samples were analyzed for bromate, and

perchlorate ions by the LC-MS/MS method (Chpater 2). ICP-MS analysis was used to

screen for transition metals ions (Co2+, Cu2+, Fe3+, Mn2+, and Ni2+). These results are

shown in Table 25.

Separately, chlorate ion was determined by iodometric titration and hypochlorite

ion by potentiometric titration with sulfite ion. Conductivity and pH measurements were

used to measure total dissolved solids, ionic strength, and pH. These results are shown in

Table 26.

Table 25. Bromate and perchlorate ions, and transition metals in bulk utility

hypochlorite ion solutions Utility

Solution ClO4

- (mg/L)

BrO3-

(mg/L)Mn

(mg/L)Fe

(mg/L)Co

(mg/L)Ni

(mg/L) Cu

(mg/L) 1 16.5 23.7 < 0.10 9.2 < 0.10 0.20 0.11 2 0.70 22.7 < 0.10 < 500 < 0.10 < 0.10 < 0.10 3 0.22 8.3 < 0.10 1.10 < 0.10 < 0.10 < 0.10 4 0.23 9.3 < 0.10 < 0.50 < 0.10 0.11 < 0.10 5 2.22 6.5 < 0.10 2.30 < 0.10 < 0.10 < 0.10

99

Table 26. Chlorate and hypochlorite ions, pH, TDS, and ionic strength in bulk utility hypochlorite ion solutions

Utility Solution

OCl-

(g/L) ClO3

- (g/L) pH

μ (mol/L)

TDS (g/L)

1 63.1 22.8 12.8 5.74 229 2 111 8.73 13.3 6.47 259 3 89.0 4.37 12.9 4.86 194 4 85.5 3.88 13.1 4.95 198 5 96.7 11.6 13.1 6.26 250

Hypochlorite ion solutions were aged at 50 ºC to determine the rate of perchlorate

ion formation and decomposition of hypochlorite ion during a 30-day incubation study.

Overlaid smoothed-line plots of perchlorate and chlorate ion formation and hypochlorite

ion decomposition in sodium hypochlorite from different utilities are shown in Figure 57.

0

20

40

60

80

100

120

0 5 10 15 20 25 30Days

OC

l- g/L

Utility 1Utility 2Utility 3Utility 4Utility 5

0

10203040506070

0 5 10 15 20 25 30Days

ClO

3-g/

L


050

100150200250300350

0 5 10 15 20 25 30

Days

mg/

L C

lO4-


Figure 57. Overlaid smoothed-line plots of (a) hypochlorite ion decomposition;

(b) chlorate ion formation; (c) perchlorate ion formation determined experimentally during incubation at 50 ºC

(a) Hypochlorite Ion Decomposition (b) Chlorate Ion Formation

(c) Perchlorate Ion Formation

100

It is evident that the hypochlorite solutions from utilities 3 and 4 behave quite

similarly (Figure 57). Thus for purposes of brevity, further discussion of hypochlorite ion

solution from utility 4 is omitted. The initial concentrations of hypochlorite and chlorate

ions, pH and temperature = 50 ºC from Table 26 were entered into Bleach 2001, to

predict hypochlorite ion decomposition and chlorate ion formation during the incubation

study. The predicted changes in concentrations of hypochlorite and chlorate ions were

compared to the measured changes, and these results are shown in Figure 58.

0

20

40

60

80

0 3 6 9 12 15 18 21 24 27 30Days

OC

l-g/

L

0102030405060

ClO

3-g/

L

Utility 1 (Experiment)Bleach 2001

0

20406080

100120

0 3 6 9 12 15 18 21 24 27 30Days

010203040506070Utility 2 (Experiment)

Bleach 2001

OC

l-g/

L

ClO

3-g/

L

020406080

100120

0 3 6 9 12 15 18 21 24 27 30

Days


Bleach 2001

OC

l-g/

L

ClO

3-g/

L

0

20406080

100120

0 3 6 9 12 15 18 21 24 27 30

Days


Bleach 2001

OC

l-g/

L

ClO

3-g/

L

Figure 58. Overlaid smoothed-line plots of hypochlorite ion decomposition and

chlorate ion formation determined experimentally in conjunction with Bleach 2001 (Error bars set at ± 10%), for solutions with (a) [OCl-]0 = 63 g/L, [ClO3

-]0 = 23 g/L;(b) [OCl-]0 = 111 g/L, [ClO3-]0 = 8.7 g/L; (c)

[OCl-]0 = 89 g/L, [ClO3-]0 = 4.4 g/L; (d) [OCl-]0 = 97 g/L, [ClO3

-]0 = 12 g/L; incubated at 50 ºC

(a) (b)

(c) (d)

101

Figure 58 shows the observed rates of hypochlorite ion decomposition, chlorate

ion formation and predicted rates of different utilities agree well within ± 10 % error bars.

As can be seen from Table 25 all these samples contain low levels of transition metals

ions, and therefore are expected to exhibit uncatalyzed decomposition of hypochlorite

ion. The Bleach 2001 predicted concentrations of hypochlorite and chlorate ions as a

function of time and the predicted second-order rate constant were used to generate the

change of perchlorate ion concentration as a function of time. An incremental increase in

perchlorate ion concentration over the initial concentration, shown in Table 25, was

calculated per time interval. The summation of incremental values over increments of

time were recorded and plotted. These results are shown in Figure 59.

0

50

100

150

200

250

0 3 6 9 12 15 18 21 24 27 30

Utility 1 (Experiment)Model

Days

ClO

4-m

g/L

0

50

100

150

200

250

300

0 3 6 9 12 15 18 21 24 27 30Days


ClO

4-m

g/L

020406080

100120140160180

0 3 6 9 12 15 18 21 24 27 30Days


ClO

4-m

g/L

0

50

100

150

200

250

300

0 3 6 9 12 15 18 21 24 27 30

Days


ClO

4-m

g/L

Figure 59. Overlaid smoothed-line plots of predicted (Error bars set at ± 10 %)

perchlorate ion formation and determined experimentally at 50 ºC in solutions with (a) [OCl-]0 = 63 g/L, [ClO3

-]0 = 23 g/L; (b) [OCl-]0 = 111 g/L, [ClO3

-]0 = 8.7 g/L(c) [OCl-]0 = 89 g/L, [ClO3-]0 = 4.4 g/L;

(d) [OCl-]0 = 97 g/L, [ClO3-]0 = 12 g/L

(a) (b)

(c) (d)

102

Figure 59 illustrates that the developed perchlorate ion model simulates the

perchlorate ion formation in bulk sodium hypochlorite solutions with pH in range of 12-

13 within ± 10 %. The error between measured and predicted concentration of

perchlorate ion increases with time. The storage of hypochlorite ion solutions over

prolonged times at elevated temperatures is highly discouraged.

5.3 Using Perchlorate Ion Formation Model to Determine Implications of Bulk Sodium Hypochlorite Solutions Storage

The goal was to determine strategies to minimize the rate of perchlorate ion

formation and thus to minimize any contribution to the treated water. However, such

strategies must also complement the current practice to slow the decomposition of

hypochlorite ion and minimize formation of chlorate ion. To assess the relative effects of

changing concentration of hypochlorite and chlorate ions, ionic strength, and

temperature, the perchlorate ion formation model was used in combination with Bleach

2001 to simulate changes in the rate of perchlorate ion formation.

In general, the “freshly” generated sodium hypochlorite solution has a lower

concentration of chlorate ion. Table 26 shows typical concentrations of hypochlorite and

chlorate ions in bulk sodium hypochlorite solutions. If the sodium hypochlorite solutions

are stored over long periods of time, the concentration of hypochlorite ion decreases and

chlorate ion increases. Thus, in general, effects of chlorate ion on the rate of perchlorate

ion are expected to be a factor during long storage periods. However, there are two key

known strategies to minimize chlorate ion formation during storage of hypochlorite ion

solutions. In practical terms, the sodium hypochlorite is diluted and/or cooled.

By diluting hypochlorite ion solutions, concentrations of both hypochlorite ion

and chlorate ion are changed. In addition dilution changes the ionic strength and pH.

Cooling the hypochlorite ion solutions reduces the second-order rate constant for

perchlorate ion formation and the rate of hypochlorite ion decomposition. Figure 60

shows overlaid smoothed-line plots of the rate of perchlorate ion formation as a function

of dilution factor or temperature. As expected, Figure 60 demonstrates that the rate of

perchlorate ion formation is reduced markedly when hypochlorite ion solution is either

103

diluted or cooled. By reducing the temperature of sodium hypochlorite solution from 35

ºC to 30 ºC, the rate of perchlorate ion formation decreases by a factor of 1.95. This

factor changes slightly as a function of temperature. For example, cooling hypochlorite

ion solutions from 50 ºC to 40 ºC reduces the rate by factor of 3.4; however, cooling a 30

ºC solution to 20 ºC results in reduction of the rate by a factor of 4.1.

Thus, cooling by as little as 5 ºC, can be an effective strategy to minimize the rate

of perchlorate ion formation approximately by a factor of 2 for any concentration of

hypochlorite ion stored below 40 ºC.

0102030405060708090

100

20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35

Temperature, ºC

% R

ate

Cooling by 5 ºC reduces the rate to 51.4% (by 1.95)

0102030405060708090

100

0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

% R

ate

Dilution Factor

1.0

At 1:2 dilution, the rate is reduced to 12.7% (by 7.9)

Figure 60. Smoothed-line plot of the rate of perchlorate ion formation as a

function of (a) temperature; (b) dilution factor. Note: Rate in solution at 2.54 M OCl-, 0.034 M ClO3

-, and μ = 7.5 M; rate at 35 ºC = 100%

(a)

(b)

104

When diluting hypochlorite ion solutions the reduction in the rate of perchlorate

ion formation is dependent on the concentration of hypochlorite and chlorate ions and

ionic strength. For example, if a hypochlorite ion solution with [OCl-]0 = 2.53 M,

[ClO3-]0 = 0.034 M, and ionic strength of 7.5 M is diluted by a factor of two, the rate of

perchlorate ion formation is decreased by a factor of 7.9 (Figure 60). Diluting by a factor

of two with 1.27 M OCl-, causes the rate of perchlorate ion formation to be reduced by a

factor of 5.6. Thus, dilution is a very effective strategy to minimize the rate of

perchlorate ion formation in concentrated sodium hypochlorite solutions. What this

suggests to a water utility is that once the bulk, concentrated sodium hypochlorite ion

solutions are delivered, to minimize the rate of perchlorate ion formation, the “fresh”

concentrated solutions should be diluted immediately.

Furthermore, diluting the concentrated sodium hypochlorite solutions is as

effective strategy as cooling but more cost-efficient. Even if storing at elevated

temperatures is necessary, dilution is a very effective strategy to minimize perchlorate ion

formation. Figure 61 (a) shows overlaid, smooth-line plots of predicted hypochlorite ion

decomposition (downward sloping curves) and perchlorate ion formation (upward

sloping curves) as a function of time at 35 ºC for hypochlorite solutions at 2.03 M OCl-

(equivalent to 13 % FAC) and at 1.02 M OCl- (equivalent to 6.5 % FAC). Dilution not

only minimizes the rate of perchlorate ion formation but also decreases the rate of

hypochlorite ion decomposition. Thus, the disinfectant solutions are stable for longer

periods of time, and when used for water-treatment, less volume will be required since

the concentration of hypochlorite ion will have degraded to a lesser extent.

This point is demonstrated by Figure 61 (b), which shows overlaid plots of μg

ClO4- per mg of FAC as a function of time (concentration of hypochlorite ion is typically

reported as mg/L FAC). Because hypochlorite ion is decomposing over time, the ratio of

μg ClO4- per mg FAC increases and thus more perchlorate ion is introduced into the

treated water per each 1 mg FAC added (1mg FAC equivalent to 0.726 mg OCl-).

105

0.000.250.500.751.001.251.501.752.002.25

0 25 50 75 100

[OC

l- ], M

0.0

20

40

60

80

100

120

Days

mg/

L C

lO4-

1.02 M OCl-2.03 M OCl-

Days

0.000.501.001.502.002.503.003.50

0 25 50 75 100

μg C

lO4- /

mg

FAC 1.02 M OCl-

2.03 M OCl-

MA MCL = 2 μg/L ClO4-


hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO4-

per mg OCl- as a function of time in solutions at 2.03 M OCl- and 1.02 M OCl- at 35 ºC

Although perchlorate ion is not regulated currently by USEPA in drinking water,

the Massachusetts Department of Environmental Protection has established a Maximum

Contaminant Level (MCL) for perchlorate ion in drinking water at 2 μg/L.103 Thus, in this

case, after 75 days of storage at 35 ºC, the concentration of perchlorate ion in treated

103 MassDEP, "Perchlorate in Public Drinking Water", 2006, Retrieved 07/21/09, from: http://www.mass.gov/dep/toxics/pchlorqa.htm

(a)

(b)

106

water would rise by 2 μg/L for every 1 mg/L FAC added (Figure 61 (b)). However, in

cases where higher doses of hypochlorite ion are required, even storing concentrated

sodium hypochlorite solutions at shorter periods may present significant contributions. In

practice, hypochlorite ion solutions are stored at room temperatures or below, and the

contribution of perchlorate ion per mg FAC is significantly lower at 25 ºC, as is

demonstrated by Figure 62.

0.000.250.500.751.001.251.501.752.002.25

0 25 50 75 1000.0

5.0

10

15

20

25

30

[OC

l- ], M

Daysm

g/L

ClO

4-

1.02 M OCl-2.03 M OCl-

0.00

0.10

0.20

0.30

0.40

0.50

0 25 50 75 100Days

μg C

lO4- /

mg

OC

l-

1.02 M OCl-2.03 M OCl-


hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO4-

per mg OCl- as a function of time in solutions at 2.03 M OCl- and 1.02 M OCl- at 25 ºC

(a)

(b)

107

In this section of the work it has been demonstrated that diluting hypochlorite ion

solutions increases the shelf-life of the disinfectant solution, while at the same time

significantly minimizes formation of perchlorate ion. Cooling diluted hypochlorite ion

solutions by as little as 5-10 ºC, further reduces the formation of perchlorate ion and also

increases the shelf-life of the disinfectant. Thus, by diluting the bulk hypochlorite ion

solutions and storing them below 30 ºC, the formation of perchlorate ion is significantly

minimized.

5.4 Application of the Perchlorate Model to OSG Sodium Hypochlorite Solutions

Because the on-site generated (OSG) sodium hypochlorite solutions (typically <1

% as FAC) are generated on demand for immediate use and thus rarely stored, these

solutions were not part of the sample matrix that was used to develop the perchlorate

predictive model. However, given the growing use of OSG sodium hypochlorite

solutions, these samples were also studied to identify the applicability of the developed

predictive model to various sodium hypochlorite solutions.

A total of 12 OSG hypochlorite ion solutions were obtained from several on-site

sodium hypochlorite generators manufacturers with different capacities, ranging from

10—2,000 pounds per day. The concentration of bromate, chlorate, and perchlorate ions

was determined by the LC-MS/MS method. ICP-MS analysis was used to determine the

concentrations of transition metals ions (Co2+, Cu2+, Fe3+, Mn2+, and Ni2+). These results

are shown in Table 27. The hypochlorite ion concentration (determined by potentiometric

titration with sulfite ion), conductivity, and pH measurements results are shown in Table

28.

108

Table 27. Transition metals, bromate, chlorate, and perchlorate ions in OSG hypochlorite ion solutions

OSG Solution

g/L ClO3-

(g/L) BrO3

-

(mg/L) ClO4

-

(μg/L) Mn

(μg/L) Fe

(μg/L) Co

(μg/L) Ni

(μg/L) Cu

(μg/L) 1 0.137 4.1 5.4 < 25 < 125 < 25 < 25 < 25 2 0.244 3.8 16 < 25 239 < 25 < 25 55 3 0.097 5.3 8.6 < 25 172 < 25 < 25 < 25 4 0.362 3.3 410 < 25 160 < 25 < 25 < 25 5 0.271 4.4 7.3 < 25 120 < 25 < 25 < 25 6 1.160 2.6 40 < 25 < 125 < 25 < 25 < 25 7 0.262 2.6 31 < 25 < 125 < 25 < 25 < 25 8 0.176 1.4 22 < 25 < 125 < 25 < 25 < 25 9 0.583 2.0 83 10 < 50 < 5.0 < 5.0 11

10 0.186 0.7 740 13 < 50 < 5.0 < 5.0 < 5.0 11 0.478 5.7 3500 < 100 < 1250 < 100 < 100 < 100 12 0.375 0.2 19 < 100 < 250 < 100 < 100 < 100

Note: Detection limits vary depending on the dilution factor used for analysis

Figure 63 shows overlaid smoothed-line plots of hypochlorite ion decomposition.

As can be seen from Table 28, the OSG hypochlorite ion solutions may have ionic

strength less than 1 mol/L and have a pH in the range of 9-10. In this pH range the

kinetics of acid-catalyzed decomposition of hypochlorite ion becomes faster and

dependent on more variables that are outside the scope of Bleach 2001. However, the

rate of hypochlorite ion decomposition is strongly dependent on the initial concentration

of hypochlorite ion and even when incubated at 50 ºC, the rate of decomposition is

relatively slow. Overlaid smoothed-line plots of perchlorate ion formation are shown in

Figure 64.

109

0

1.52.53.54.55.56.57.5

0 5 10 15 20 25 30Days

OC

l-g/

L

OSG 1OSG 2OSG 3OSG 4OSG 5OSG 6

0.5

00.51.52.53.54.55.56.57.5

0 5 10 15 20 25 30

Days

OSG 7OSG 8OSG 9OSG 10OSG 11OSG 12O

Cl-

g/L

Figure 63. Overlaid smoothed-line plots of hypochlorite ion decomposition (a)

OSG solutions 1-6; (b) OSG solutions 7-12, 50 ºC Table 28. Concentration of hypochlorite ion, pH, TDS, and ionic strength (μ) in

OSG hypochlorite ion solutions OSG

Solution μ

(mol/L) TDS (g/L) pH

OCl- (g/L) % FAC

1 0.669 26.7 9.4 7.1 0.95 2 0.589 23.5 9.3 5.8 0.78 3 0.945 37.8 9.1 4.9 0.65 4 0.526 21.1 9.2 5.0 0.67 5 0.682 27.3 9.3 7.4 0.99 6 0.502 20.1 8.8 3.3 0.44 7 1.153 46.1 9.1 5.8 0.77 8 0.546 21.8 9.4 3.8 0.51 9 0.923 36.9 9.5 5.2 0.69 10 0.265 10.6 9. 6 2.6 0.35 11 0.828 33.1 9.4 4.9 0.66 12 0.882 35.3 9.4 6.3 0.84

(a)

(b)

110

050

100150200250300350

0 5 10 15 20 25 30Days

0

100

200

300

400

500

600

ClO

4-μg

/L

OSG 1OSG 2

OSG 3OSG 4

OSG 5OSG 6

ClO

4-μg

/L (O

SG 4

)

050

100150200250300350

0 5 10 15 20 25 30Days

0500100015002000250030003500400045005000

ClO

4-μg

/L

ClO

4-μg

/L (O

SG 1

0, 1

2)OSG 7OSG 8

OSG 9 OSG 10OSG 11 OSG 12

Figure 64. Overlaid smoothed-line plots of perchlorate ion formation (a) OSG

solutions 1-6; (b) OSG solutions 7-12, 50 ºC

As can be seen from Figure 64, the rate of perchlorate ion formation varies by

OSG manufacturer. An average rate over the duration of incubation period was

calculated and used to determine experimental values for the second-order rate constant,

using Equation 36. By using the measured values of ionic strength and temperature, the

predictive model was used to determine the second-order rate constants. Table 29 shows

the experimental and predicted values of the second-order rate constants of perchlorate

ion formation.

(a)

(b)

111

Table 29. Second-order rate constants of perchlorate ion formation in OSG hypochlorite ion solutions, experiment vs. model (k2obs= experimental k2, k2cal = predicted k2, in units of L·mol-1·day-1)

OSG k2obs k2cal k2obs/k2cal 1 302 89.9 3.4 2 143 88.7 1.6 3 406 94.6 4.3 4 109 87.7 1.2 5 203 90.2 2.3 6 24.6 87.3 0.3 7 202 98.2 2.1 8 95.7 88.0 1.1 9 117 94.2 1.2 10 234 83.6 2.8 11 259 92.6 2.8 12 169 93.5 1.8

The effects of pH were discussed in Chapter 3, and the results shown in Table 29

confirm the fact that the rate of perchlorate ion formation is faster in the OSG

hypochlorite ion solutions than in bulk sodium hypochlorite. There is no observable

correlation between the experimentally-observed values of the second-order rate constant

and pH or ionic strength of these solutions; however during this work, it was clearly

shown that the formation of perchlorate is catalyzed at lower pH values. This is an

indication that the formation of perchlorate ion in this type of sample matrix (low pH

values and low ionic strength) is dependent on more variables than the bulk, concentrated

hypochlorite ion solutions. Therefore, although the model may not accurately predict the

formation of perchlorate ion in OSG hypochlorite ion solutions, it does provide

groundwork for additional research.

However, it is clear from these results that perchlorate ion does form in OSG

hypochlorite ion solutions. Thus, to minimize the formation of perchlorate ion, it is

recommended not to store the OSG hypochlorite ion solutions at elevated temperatures or

for more than one to two days, which is typically the practice.

112

5.5 Application of the Perchlorate Model to Calcium Hypochlorite Solutions

As an alternative to sodium hypochlorite, calcium hypochlorite is used for

disinfection of water. Calcium hypochlorite, typically available as a hydrated salt, is

stored as solid and is dissolved upon use. Recently it has been reported that the rate of

calcium hypochlorite decomposition is faster than that of sodium hypochlorite in

solutions at elevated temperatures.104 It also has been reported that perchlorate ion is not

a common contaminant in calcium hypochlorite.105

A sample of calcium hypochlorite, 60-80 % as Ca(OCl)2, was obtained from

Arch Chemicals (Norwalk, CT). Two stock solutions containing approximately 22 g/L

and 44 g/L OCl- (~3 % and 6 % FAC) were prepared. Separate aliquots were stored for

two weeks at 60 ºC and four weeks at 50 ºC. Figure 65 shows smoothed-line plots of

hypochlorite ion decomposition (downward sloping curves) and chlorate ion formation

(upward sloping curves).

Trace amounts of perchlorate ion were detected in both stock calcium

hypochlorite solutions by the optimized LC-MS/MS method described in Chapter 2. The

average perchlorate ion concentration of 624 ± 29 μg/kg in calcium hypochlorite was

determined by the analysis of duplicate samples. The concentration of perchlorate ion

increased in calcium hypochlorite solutions during incubation at 50 ºC and 60 ºC and is

shown in Figure 66.

104 Su, Y. S., Morrison, D. T. and Ogle, R. A. "Chemical kinetics of calcium hypochlorite decomposition in aqueous solutions" J. Chem. Health Safety 2009, 16, 21-25. 105 Greiner, Mclellan, Bennet and Ewing "Occurrence of perchlorate in sodium hypochlorite".

113

05

101520253035404550

0 4 8 12 16 20 24 28Days

02468101214161820

OC

l-g/

L

ClO

3-g/

L

[OCl-]0= 44 g/L; [ClO3-]0=0.8 g/L

[OCl-]0= 23 g/L; [ClO3-]0=0.4 g/L

05

101520253035404550

0 2 4 6 8 10 12 14Days

OC

l-g/

L

02468101214161820

ClO

3-g/

L

[OCl-]0= 45 g/L; [ClO3-]0=0.8 g/L

[OCl-]0= 23 g/L; [ClO3-]0=0.4 g/L

Figure 65. Smoothed-line plots of hypochlorite ion decomposition and chlorate

ion formation in calcium hypochlorite solutions (a) incubated at 50 ºC; (b) incubated at 60 ºC

0123456789

10

0 4 8 12 16 20 24 28Days

ClO

4-m

g/L

[OCl-]0= 44 g/L; [ClO3-]0=0.8 g/L

[OCl-]0= 23 g/L; [ClO3-]0=0.4 g/L

012345678

0 2 4 6 8 10 12 14Days

ClO

4-m

g/L

[OCl-]0= 45 g/L; [ClO3-]0=0.8 g/L

[OCl-]0= 23 g/L; [ClO3-]0=0.4 g/L

Figure 66. Smoothed-line plots of perchlorate ion formation in calcium

hypochlorite solutions (a) incubated at 50 ºC; (b) incubated at 60 ºC

Similar to sodium hypochlorite, the concentration of perchlorate ion increases in

stored calcium hypochlorite solutions and is dependent on concentration of hypochlorite

and chlorate ions, ionic strength, and temperature. The rate of perchlorate ion formation

at 60 ºC was measured at 0.76 μmol/L per day, and at 50 ºC, the rate was 0.20 μmol/L per

day in solution at 23 g/L OCl- and 0.4 g/L ClO3-. Thus, if the temperature is increased by

10 ºC, the rate increased by a factor of 3.8, which is higher than the predicted factor of

3.2 for sodium hypochlorite solutions. Table 30 shows the calculated and measured

second-order rate constants for different calcium hypochlorite solutions.

(a) 50 ºC (b) 60 ºC

(a) 50 ºC (b) 60 ºC

114

Table 30. Second-order rate constants of perchlorate ion formation in calcium

hypochlorite solutions, experiment vs. model (k2obs= experiment k2, k2cal = k2

predicted in units of L·mol-1·day-1)

T, ºC

[OCl-] (mol/L)

[ClO3-]

(mol/L) kobs kcal k2obs / k2cal pH

Ionic strength

50 0.446 0.005 96 91.2 1.0 11.2 0.75 50 0.859 0.010 133 102 1.3 11.5 1.38 60 0.443 0.005 332 291 1.1 11.4 0.75 60 0.874 0.010 594 326 1.8 11.5 1.38

The results shown in Table 27 indicate that the observed second-order rate

constants are higher than predicted (as noted in k2obs / k2cal column). This may be

attributed to the difference in the mean activity of the 1:1 electrolyte solution of sodium

hypochlorite (NaOCl) and 1:2 electrolyte solution of calcium hypochlorite106 (Ca(OCl)2).

Similar to OSG hypochlorite ion solutions, it is recommended not to store the

calcium hypochlorite solutions for more than several days. By using diluted solutions of

calcium hypochlorite and/or cooling while storing, formation of perchlorate ion is

significantly reduced. A survey of calcium hypochlorite from different manufacturers

may reveal additional information that may be useful for comparison to perchlorate ion

formation in sodium hypochlorite solutions.

5.6 Potential Contribution of Perchlorate Ion to Drinking Water from Various

Hypochlorite Ion Solutions

As was noted earlier, the participating utilities provided water samples along with

the bulk sodium hypochlorite solutions. Distribution samples consisted of “finished”

(treated) water at different residence times, shown in Table 31, and are denoted as

Residence (Res.) Time A and Res. Time B. Samples of raw, finished, and distribution

system waters were analyzed for perchlorate, bromate, and chlorate ions by the LC-

MS/MS method (Chapter 2), and results are shown in Tables 32-34.

106 Robinson and Stokes Electrolyte Solutions.

115

Table 31. Residence time of the sampled distribution waters

Utility Res. Time (A) (hrs) Res. Time (B) (hrs) 1 36 72 2 72 216 3 36 72 4 12 24 5 100 150

Table 32. Perchlorate ion in raw, finished, and distribution waters ClO4

- Raw Finished Time A Time B

Utility (μg/L) (μg/L) (μg/L) (μg/L) 1 < 0.5 3.6 < 0.5 3.10 2 < 0.5 < 0.5 < 0.5 < 0.5 3 < 0.5 < 0.5 < 0.5 < 0.5 4 < 0.5 < 0.5 < 0.5 < 0.5 5 1.6 1.2 1.2 0.9

Table 33. Bromate ion in raw, finished, and distribution waters BrO3


Utility (μg/L) (μg/L) (μg/L) (μg/L) 1 < 0.5 0.5 0.8 2.9 2 < 0.5 < 0.5 < 0.5 < 0.5 3 1.3 1.4 2.1 2.2 4 < 0.5 0.9 0.8 0.9 5 < 0.5 2.6 5.9 3.2

Table 34. Chlorate ion in raw, finished, and distribution waters ClO3


Utility (μg/L) (μg/L) (μg/L) (μg/L) 1 14.0 583 590 1200 2 5.0 19.0 46.0 45.0 3 130 198 < 3.0 < 3.0 4 < 3.0 129 126 133 5 < 3.0 788 1590 823

Results in Table 32 indicate that the levels of perchlorate ion do not increase over

time in the treated waters. None of the water samples contained perchlorate ion

116

concentration above the USEPA established interim health advisory level107 of 15 μg/L.

Bromate ion concentration was measured below an MCL of 10 μg/L in all water samples

(Table 33). However, chlorate ion was present in several samples above the guideline

MCL of 700 µg/L (Table 34), set by the World Health Organization.108 This suggests

that addition of contaminants such as chlorate and bromate ions during disinfection

process may be more of an immediate concern to the drinking water industry, than

potential addition of perchlorate ion (given that a water utility practices strategies to

minimize decomposition of hypochlorite ion and formation of chlorate ion).

To evaluate any potential contributions of perchlorate, bromate, and chlorate ions

from the hypochlorite ion solutions used to disinfect the water, a ratio of concentration of

contaminant per mg FAC was calculated and is presented in Table 35.

The results shown in Tables 26 and 35 support a key point that the bulk

hypochlorite ion solutions that are stored for long periods have higher concentrations of

contaminants introduced into the treated water. For example, the solution obtained from

Utility 1 contained 63.1 g/L OCl- and 22.8 g/L ClO3-, which indicates that this solution

has been decomposing. Solution from Utility 3 contained 89.0 g/L OCl- and 4.37 g/L

ClO3-, which indicates that this solution is “fresher” (Table 26). The ratio of contaminant

per mg FAC is much higher for the solution from Utility 1. Perchlorate ion concentration

in treated water from Utility 1 was significantly higher than in the sample obtained from

Utility 3. Thus, based on the limited data, it can be concluded that using “fresh” solutions

of hypochlorite ion would introduce lower concentrations of contaminants. This leads to

the conclusion that concentrated sodium hypochlorite solutions should be diluted at the

time of delivery and cooled in order to minimize decomposition of the hypochlorite ion

and the concomitant formation of chlorate and perchlorate ions.

The OSG sodium hypochlorite solutions are typically generated and used based

on demand and are not stored for more than one or two days. However, these solutions

vary significantly in the ratio of different contaminant per mg FAC. For example, OSG

solution 12, at 8.7 g/L FAC, had a considerably lower ratio of contaminants per mg FAC.

This emphasizes that there are multiple variables involved during hypochlorite ion

107 USEPA "Interim Drinking Water Health Advisory for Perchlorate" 108 WHO Guidelines for Drinking Water Quality; World Health Organization: Singapore, 2006, http://www.who.int/water_sanitation_health/dwq/gdwq3rd_add1.pdf

117

generation. Salt purity and bromide ion levels in source water, for example, may be a

significant factor in the levels of bromate ion in the generated hypochlorite ion solution.

Thus, the levels of perchlorate, bromate, and chlorate ions should be routinely

monitored and suggested OSG manufacturer’s salt specification and maintenance

guidelines considered.

Table 35. Contributions of perchlorate, bromate, and chlorate ions per mg FAC

in various hypochlorite ion solutions

Bulk Sodium Hypochlorite

FAC (g/L )

ng ClO4-

per mg FAC

ng BrO3-

per mg FAC

μg ClO3-

per mg FAC

1 87.0 190 272 262 2 153 4.58 148 57 3 123 1.79 67.7 36 4 118 1.95 78.9 33 5 133 16.7 48.8 87

OSG Sodium Hypochlorite

FAC (g/L )

ng ClO4-

per mg FAC

ng BrO3-

per mg FAC

μg ClO3-

per mg FAC

1 9.7 0.55 421 14 2 8.0 2.00 475 30 3 6.8 1.27 782 14 4 6.9 59.9 482 53 5 10.2 0.72 432 27 6 4.5 8.9 576 257 7 8.0 3.85 323 33 8 5.2 4.25 271 34 9 7.2 11.6 280 81 10 3.6 208 199 52 11 6.8 515 839 70 12 8.7 2.18 17.2 43

Calcium Hypochlorite

FAC (g/L )

ng ClO4-

per mg FAC

ng BrO3-

per mg FAC

μg ClO3-

per mg FAC

1 31.5 0.84 77.2 13 2 61.5 0.86 79.8 13

As an alternative, small drinking-water utilities may consider the use of calcium

hypochlorite for disinfection treatments. Calcium hypochlorite contained the lowest ng

ClO4- per mg FAC, and one of lowest ng BrO3

- per mg FAC and μg ClO3- per mg FAC,

as can be seen in Table 35.

118

5.7 Conclusions

The use of Equation 50 to determine the value of second-order rate constant as a

function of the ionic strength of solution at specific temperatures has been validated. A

difference of less than ± 10 % between the experimentally-measured decomposition of

bulk hypochlorite ion and that predicted by Bleach 2001 was demonstrated. Thus, by

using Equation 48 and predicted concentrations of hypochlorite and chlorate ions, a

simple predictive model is provided for perchlorate ion formation in bulk sodium

hypochlorite:

)eeT10log(2.084)0.0788()log(k R106

RT1.01x10

102

5 −−

××××+= μ (48)

• Application of the predictive model on several bulk sodium hypochlorite solutions

obtained from different water utilities demonstrated an accurate estimate of

perchlorate ion formation in the bulk sodium hypochlorite solutions.

• Perchlorate ion concentration increased in OSG sodium hypochlorite and calcium

hypochlorite solutions, exhibiting dependence on more variables than the bulk

sodium hypochlorite solutions. This is due to differences in pH and ionic strength.

• The results of a survey of raw and treated waters demonstrated that currently

perchlorate ion is present at trace concentrations (< 5.0 μg/L). However, the

results revealed that the levels of chlorate ion are considerably higher than the

WHO MCL109 of 700 μg/L. Thus, strategies to minimize formation of chlorate

ion in hypochlorite ion solutions should be considered by the utilities.

• It is evident that the use of stored concentrated hypochlorite ion solutions should

be avoided. Significant reductions in the rate of hypochlorite ion decomposition

and in the rates of chlorate ion and perchlorate ion formation can be achieved by

dilution and cooling of hypochlorite ion solutions.

109 WHO Guidelines for Drinking Water Quality

119

CHAPTER 6. CONCLUSIONS

6.1 Summary

A robust, sensitive LC-MS/MS method was developed and used for the

determination of perchlorate, bromate, and chlorate ions. Potentiometric titration with

sulfite ion was used for the determination of hypochlorite ion, and an iodometric titration

method was used for the determination of chlorate ion in bulk hypochlorite ion solutions.

Proper storage and sample preservation conditions were established.

The results of this study indicate that perchlorate ion formation in sodium

hypochlorite solutions is dependent on several factors: (1) The concentration of

hypochlorite and chlorate ions directly impact perchlorate ion formation; (2) The

presence of transition metal ions, chlorite ion or bromide ion indirectly impact

perchlorate ion formation, by reacting with hypochlorite ion; (3) The presence of noble

metal ions or bromate ion has no observable effect on perchlorate formation; (4) Higher

ionic strength enhances perchlorate ion formation; (5) pH effects in concentrated

hypochlorite ion solutions are more dominant in hypochlorite ion decomposition (faster

kinetics) than in perchlorate ion formation; however, in more dilute (i.e. more stable)

hypochlorite ion solutions, perchlorate ion formation also appears to be acid-catalyzed.

The rate of perchlorate ion formation was found to be first-order with respect to

hypochlorite and chlorate ions and second-order overall. The rate constant was found to

be strongly dependent on the ionic strength and a reduced expression to account for

effects of the ionic strength and temperature was validated. As a result, a perchlorate ion

formation predictive model is provided. Application of the predictive model with Bleach

2001 on several bulk sodium hypochlorite solutions obtained from different water

utilities, demonstrated an accurate approximation of perchlorate ion formation in bulk

sodium hypochlorite solutions.

Perchlorate ion concentration increased in the OSG sodium hypochlorite and

calcium hypochlorite solutions, exhibiting dependence on more variables than in the bulk

sodium hypochlorite solutions over time, due to differences in pH and the ionic strength.

The results of a survey of raw and treated waters demonstrated that currently

perchlorate ion is present at trace concentrations (< 5.0 μg/L). However, the results

120

revealed that the levels of the chlorate ion are considerably higher than the WHO MCL110

of 700 μg/L. Thus, strategies to minimize formation of the chlorate ion in hypochlorite

ion solutions should be considered by utilities.

It is evident that the use of stored concentrated hypochlorite ion solutions should

be avoided. Significant reductions in the rate of hypochlorite ion decomposition and in

the rates of chlorate ion and perchlorate ion formation can be achieved by dilution and

cooling of the stored hypochlorite ion solutions.

6.2 Recommendations to Water Utilities

Based on the results of this study the following recommendations to water utilities

are presented in light of the potential regulation of perchlorate ion in the drinking water:

i. Dilute Bulk Sodium Hypochlorite Solutions Upon Delivery

By diluting a hypochlorite ion solution at 2 M OCl- by a factor of 2, the

rate of perchlorate ion formation decreases by a factor of 7. If the same

solution is diluted by a factor of ten, the rate of perchlorate ion formation

decreases by a factor of 266.

ii. Store the Hypochlorite Ion Solutions At Lower Temperatures

By cooling sodium hypochlorite solutions by 5 ºC, the rate of perchlorate

ion formation is reduced approximately by a factor of 2. Thus, seasonal

changes in the temperature should be taken into consideration when

storing hypochlorite ion solutions.

iii. Maintain the pH of Bulk Hypochlorite Ion Solutions above pH 11 and

Below pH 13 After Dilution

By minimizing the decomposition of hypochlorite ion, a lower ratio of the

contaminants per mg FAC is achieved. Chlorate and perchlorate ions will

increase during prolonged storage. By keeping the pH in the range of

11-13, the decomposition of hypochlorite ion is minimized. 110 WHO Guidelines for Drinking Water Quality

121

iv. Require Removal of Transition Metal Ions From the Bulk Sodium

Hypochlorite Manufacturer

The presence of transition metal ions catalyzes decomposition of

hypochlorite ion thus increasing the ratio of the contaminant per mg FAC.

v. Do Not Store OSG and Calcium Hypochlorite Solutions

OSG sodium hypochlorite solutions are at pH ~9.5 and the rate of

perchlorate ion formation is enhanced. However, the rate of perchlorate

ion formation is still dependent on concentration of both hypochlorite and

chlorate ions. Thus, the on-site generators should be maintained and

optimized to reduce loss of hypochlorite ion to form chlorate ion.

Therefore, OSG solutions should be generated for immediate use and

should not be stored for more than 1-2 days. Perchlorate ion will also

form in calcium hypochlorite solutions and thus only fresh solutions

should be used.

vi. Use a Higher Purity Salt for OSG Sodium Hypochlorite, to Minimize

Conversion of Bromide Ion to Bromate Ion

Bromide ion reacts rapidly with hypochlorite ion to produce hypobromite

ion, which decomposes rapidly to produce bromate ion. Thus, higher

purity salts containing lower amounts of bromide ion will results in less

bromate ion formed during the on-site generation, and thus contribution of

the bromate ion to the drinking water will be significantly minimized.

122

APPENDIX 1. DETECTION OF OZONE GAS BY GOLD NANOISLANDS

A1.1 Introduction

In a previous paper this group reported that aqueous solutions of ozone produced a

shift in the surface plasmon resonance from colloidal gold nanoparticles.111 In most ozone

methods, ozone irreversibly oxidizes the reagents thereby making a sensor based on that

reagent impractical. The surprising finding was that the shift in the surface plasmon

resonance from the colloidal gold nanoparticles reversed to the original wavelength upon

depeletion of ozone. This is surprising since most reports of analytical utilization of

surface plasmons claimed irreversibility since the plasmonic shift was due to irreversible

aggregation of the gold nanoparticles. It was observed that the colloidal gold nanoparticles

were able to reproducibly cycle between the surface plasmon wavelengths as ozone was

introduced and removed. Since the desired outcome was a reversible ozone sensor, the

colloidal solution had to be replaced by a solid state surface.

In that initial report gold nanoislands were made using vapor deposition and were

tested for gaseous ozone response. As was the case in the aqueous system, the surface

plasmon resonance of the gold nanoislands exhibited a cycling effect when ozone was

introduced and removed. The difficulty with this approach was that the process of using

vapor deposition followed by a heat treatment was not very reproducible. The biggest issue

was controlling the size of the gold nanoislands. Described in this paper is a new process

to make more reproducible gold nanoislands and the analytical characterization of this new

material in terms of response to ozone.

The objective is to develop an ozone sensor for the drinking water industry and

other industries that electrolytically produce ozone. Tighter control of the generation

system will produce significant energy cost savings while maximizing the efficacy of the

ozone treatment. Up to now this task has been complicated by the complex nature of ozone

decomposition pathways112 and irreversible oxidation of the reagents/sensing element.113

111 Puckett, S. D., Heuser, J. A., Keith, J. D., Spendel, W. U. and Pacey, G. E. "Interaction of ozone with gold nanoparticles" Talanta 2005, 66, 1242-1246. 112 Tomiyasu, H., Fukutomi, H. and Gordon, G. "Kinetics and mechanism of ozone decomposition in basic aqueous solution" Inorg. Chem. 1985, 24, 2962-2966. 113 Gordon, Cooper, Rice, and Pacey Disinfectant Residual Measurement Methods.

123

Electrochemically-based measurements have also been problematic due to poor sensitivity

and electrode fouling.114

This paper discusses the characterization of the gold nanoisland surface, the gold

nanoisland interaction with ozone, the desirable performance characteristic(s) of the

sensor, the optimal size for gold nanoisland, the current sensor configuration, and the

sensor performance characteristics.

A1.2 Experimental

Ozone generation was performed using a water-cooled Ozone Research

Equipment Corporation Model O3B9-0 generator. The cooling water and fan are turned

on first. Oxygen pressure from the oxygen cylinder is reduced down to 6 psi by a

regulator. The generator is turned on and the electric current set to 1.3 amps. This

setting should produce approximately 9 grams per hour of ozone. Ozone concentration

was determined by measuring absorbance at 254 nm on HP 8453 Photodiode UV-Vis.

Molar absorptivity of 3300 M-1 cm-1 was used to calculate the concentration of ozone. A

10 cm quartz sample cuvette was used to increase sensitivity of UV-Vis method. Perkin

Elmer, Lambda 950 Reflectance Spectrophotometer, employing an integrating sphere

was used to collect absorbance spectra of gold nanoislands. Varian’s Cary 50

spectrophotometer was used to collect UV-Vis spectra of gold nanoislands sputtered onto

transparent substrates, such as polished quartz and indium tin oxide coated glass plates.

Previously,115 the gold nanoislands were produced using Physical Vapor

Deposition (PVD) onto glass or Indium Tin Oxide coated glass. The current work uses a

sputtering procedure to produce gold nanoislands. This greatly reduces synthesis time

and it also provides very good control of sputtering parameters, which improves

reproducibility. Gold nanoislands were generated using Anatech Hummer V Sputter

Coater. Typical settings for the sputter coater: 70-80mtorr vacuum, High Voltage set to

generate current above 10 mA and coating time: 1 min-5min. The next step is to anneal

the gold coating at 375°C for 15 min. This temperature was selected to reduce the time

114 Gordon, Cooper, Rice, and Pacey Disinfectant Residual Measurement Methods. 115 Puckett, Heuser, Keith, Spendel, and Pacey "Interaction of ozone with gold nanoparticles".

124

to produce nanoislands. Lower temperatures require longer annealing time but can be

used as well.

The synthezid nanoislands was characterized by an Agilent Technologies 5500

Atomic Force Microscope and Carl Zeiss Supra Scanning Electron Microscope.

A1.3 Results and Discussion

Upon making the gold nanoisland surface, AFM and SEM images of the surface

were obtained. Figure A1-1 shows a typical SEM image of gold nanoislands on polished

aluminum substrate. Figure A1-2 shows the AFM image of AFM images of 25 nm (a);

and 14 nm (b) gold nanoislands on quartz substrate. Unlike the prior work, the

nanoislands are evenly distributed across the surface.

Figure A1-1. SEM images of typical gold nanoislands produced by sputtering process on a polished aluminum substrate

125

Figure A1-2. AFM images of 25 nm (a); and 14 nm (b) gold nanoislands on quartz substrate

As can be seen from Figure A1-3, gold nanoislands respond to the presence of

ozone by producing a red-shift in the surface plasmon resonance. The surface plasmon

resonance peak for the gold nanoislands shifts from 520 nm to 540 nm upon exposure to

ozone. The presence and concentration of ozone was monitored by measuring the ozone

absorbance peak at 254 nm. Upon decomposition of ozone from the sample solutions, the

(a)

(b)

126

surface plasmon resonance peak for the gold nanoislands completely return to the original

position at 520 nm.

The top line in Figure A1-3 shows the spectrum for a thin film of gold that produces

a broad absorbance after 600nm and no surface plasmon resonance. In previous studies

surface spectroscopy indicated that the ozone decomposed to oxygen which covered the

surface.116 Formation of gold oxides and chemisorbed oxygen layers was also reported117.

The oxygen layer could be removed by heating the surface to 600 K and that this

desorption process exhibited first-order kinetics.118

Wavelength, nm

200 400 600 800 1000

Abs

orba

nce

0.00

0.05

0.10

0.15

0.20

0.25Sputtered gold thin filmGold nanoislands Gold nanoislands exposed to OzoneGold nanoislands reversed

Figure A1-3. Overlaid UV-Vis spectra of : dashed line-gold thin film on quarts, solid line-gold nanoislands, dot-dashed line gold nanoislands exposed to ozone gas, dotted line-gold nanoislands reversed by annealing at 375 ºC for 15 min

116 Saliba, N., Parker, D. H. and Koel, B. E. "Adsorption of oxygen on Au(111) by exposure to ozone" Surf. Sci. 1998, 410, 270-282. 117 Krozer, A. and Rodahl, M. "X-ray photoemission spectroscopy study of UV/ozone oxidation of Au under ultrahigh vacuum conditions" J. Vac. Sc. Technol. A. 1997, 15, 1704-1709. 118 Saliba, Parker and Koel "Adsorption of oxygen on Au(111) by exposure to ozone".

127

Fortunately, gold nanoislands do not need high temperatures to shift the surface

plasmon resonance back to 520 nm. This suggests that unlike with the thin fold films,

where the ozone decomposed to oxygen which was adsorbed, the ozone is directly

influencing the surface plasmon resonance and when it decomposes, the oxygen that is

produced does not interact with the gold nanoislands. Studies of pure oxygen confirmed

this finding. Exposure of gold nanoislands to oxygen gas does not cause a surface-plasmon

resonance to shift.

In addition shifts in surface plasmon resonance has been explained by aggregation.

Using particle size analysis it has been shown that the size increases (doubles or more)

when the shift is observed. But in the case of ozone particle size analysis shows that the

size has not changed when the shift is observed. The ozone on the particle alone is creating

the shift. This also explains why the ozone produced shift is reversible. Aggregated based

shifts do not reverse since breaking the aggregation is not probable. For the ozone based

shift, the ozone only has to decompose leaving the same size of the nanoparticle that

existed before the ozone exposure.

The desirable performance characteristic(s) of an ozone sensor have been

previously described.119 The ideal method/technique would be specific for only ozone

itself and not respond to its decomposition products. It needs to be selectivity over possible

interferences by at least 500. A detection limit of 0.01 mg/L for ozone with a precision of

at least 0.1% and an accuracy of at least 0.5% is needed. The system must be easily,

absolutely calibrated. A working linear range of five orders of magnitude is desirable but

not mandatory for all situations. The system can not require dilution of the water or gas

sample to mask or minimize potential interferences. The system must work equally well

(and satisfy the above requirements) in batch and automated modes of operation. The

system must be operable by a technician and not require specialized skills or unusual or

ultra-complex instruments. The measurement should be made in less than one minute. The

whole process must be relatively cost effective.

The response of the gold nanoislands sensor to ozone was investigated by

exposing it to different concentrations. Figure A1-4 shows the spectra exhibiting the

amount of red-shift, in nm, with increasing ozone concentration. Figure A1-5 is a

119 Gordon, Cooper, Rice, and Pacey Disinfectant Residual Measurement Methods.

128

calibration plot of spectral shift (nm) of the gold nanoislands surface plasmon resonance

peak versus ozone concentration. The calibration plot shows that with the current design

the detection of gaseous ozone at concentration of less than 20 ppb (µg/L). Starting with

the 20 ppb data point, a curve with an equation of y=6.8ln(x)-15.19 produced a

correlation coefficient of 0.966.

Wavelength (nm)

460 480 500 520 540 560 580 600

Abso

rban

ce

0.035

0.040

0.045

0.050

0.055

0.060

0.065

0.070

0.075

Figure A1-4. Overlaid UV-Vis spectra of 25 nm gold nanoislands with surface Plasmon absorbance max at 520 nm exposed to concentrations of ozone, increased in increments form 20.9 μg/L to 166.1 μg/L. Ozone causes a red-shift in the surface-plasmon absorbance max

129

Ozone Concentration, ppb (ug/L)

0 20 40 60 80 100 120 140 160 180

Shift

, nm

02468

10121416182022

Figure A1-5. Shift of the 25 nm gold nanoislands surface-plasmon max (520 nm) as a function of ozone concentration, logarithmic fit gives an equation of y=6.8ln(x)-15.19 produced a correlation coefficient of 0.966

So far the majority of this work is performed with 10-30 nm gold nanoislands.

However; based on this group’s aqueous gold nanoparticle work, we hypothesized that the

size and density of gold nanoislands affects the response performance of the sensor. The

current work shows evidence that different size and density of gold nanoislands may in fact

affect the performance of sensor. As evidence of this effect, Figure A1-6 shows a plot of

surface-plasmon max shift as a function of ozone concentration of gold nanoislands with

absorbance max of 532 nm. This sample of gold nanoislands has absorbance max at a

longer wavelength than the 25 nm gold nanoislands with absorbance max of 520 nm.

Absorbance max of the surface plasmon suggests that the nanoislands shown in Figure A1-

6, are of a larger size. Figure A1-6 demonstrates that different sizes of gold nanoislands

are saturated by ozone at different concentrations. This may also suggest that the smaller

particles are more densely packed and thus interparticle distance is smaller. The

interparticle distance is known to affect the surface plasmon shift and thus the amount of

ozone that is adsorbed between the particles may in fact also be dependent on the spacing

130

of the particles. It would also make sense to extend the range of nanoislands to sizes

between 1 and 100 nm. It is also possible that an array of different sizes of gold

nanoislands can be employed synergistically to determine ozone present over wide range of

concentrations.

Ozone Concentration, ppb (ug/L)

0 80 160 240 320 400 480 560 640 720 800 880 960 1040

Shift

, nm

10

12

14

16

18

20

Figure A1-6. Surface plasmon’s shifts of gold nanoislands with absobance max at 532 nm as a function of ozone concentration

Future sensor designs will place gold nanomaterials in a substrate material that is

resistant to ozone, cost effective, and interference free to the detection scheme. The

substrate does not need to be transparent in the UV-Vis region of electromagnetic

spectrum, since a reflective mode can be used to collect absorbance spectra of the gold

nanoislands. The substrate should also provide open access to the surface of gold

nanoislands. Lastly, our current method to reverse the spectral shift of gold nanoislands

properties is performed by heating the nanoislands. The reversal time is dependent on the

temperature used. At 375 ºC the optical shift is reversed by 10-15 minutes. At 250 ºC 20-

25 minutes may be required to fully reverse the optical shift of gold nanoislands caused

by ozone. Heating the gold nanoislands drives the ozone off and thus reverses the optical

131

shift. No interferences were observed for reversing the gold nanoislands surface Plasmon

by annealing process. Several samples were exposed for significantly longer times and

found that the surface plasmon peak does disappear if the gold nanoislands are heated for

longer periods of time. This is most likely due to forming a gold oxide layer which

inactivates the nanoislands.

A1.4. Conclusions

It has been demonstrated that fabrication of fast and cheap sensor material that

can be used for detection of ozone is possible. Surface Plasmon resonance shift of gold

nanoislands is reversible. We have tested a consistent reversibility has been observed. A

detection limit of 15 ppb was achieved.

132

APPENDIX 2. ELECTROCHEMICALLY ASSISTED PROCESSING OF

ORGANICALLY MODIFIED, PERPENDICULARLY

ORIENTED MESOPOROUS SILICA FILMS WITH

FLUORESCENT FUNCTIONALITY

A2.1. Introduction

Functionalized, mesostructured thin films are of interest in a number of fields,

including the development of optical and electrochemical sensing devices. A key to

optimizing the performance of such sensors is to provide for facile mass transport into

and through the film, which suggests fabrication with a controlled pore structure. In this

regard, the least tortuous path is achieved with pores normal to a substrate such as an

electrode. Ordered mesoporous supports have been achieved in the past by such

approaches as optical lithography,120 ion-track etching,121 electrochemical etching of

Si122, and anodization of Al123. However, these techniques, which impart porosity to an

existing film, suffer from problems such as restricted pore density or disordered pore

structures. Others have identified synthetic approaches to produce structurally controlled

mesoporous thin films with the desired orientation.124,125,126,127 Our objective is to

directly synthesize mesoporous thin films that have cylindrical pores normal to the

substrate and that are functionalized with sensing centers, such as fluorophores.

Early studies of surfactant-templated materials used a liquid-crystal template

mechanism in which aluminosilicate gels were calcined in the presence of surfactant- 120 Choi, Y.-K., King, T.-J. and Hu, C. "Nanoscale CMOS spacer FinFET for the terabit era" Electron. Devic. Lett. 2002, 23, 25-27. 121 Metz, S., Trautmann, C., Bertsch, A. and Renaud, P. "Polyimide microfluidic devices with integrated nanoporous filtration areas manufactured by micromachining and ion track technology" J. Micromech. Microeng. 2004, 14, 324-331. 122 Sato, H., Homma, T. "Fabrication of high-aspect-ratio arrayed structures using Si electrochemical etching" Sci. Technol. Adv. Mat. 2006, 7, 468-474. 123 Xu, C. X., Zhang, X. S. and Sun, X. W. "Preparation of Porous Alumina by Anodization" J. Metastab. Nanocryst. 2005, 23, 75-78. 124 Sanchez, C., Julian, B., Belleville, P. and Popall, M. J. Mater. Chem. 2005, 15, 3559. 125 Scott, B. J., Wirnsberger, G. and Stucky, G. D. "Mesoporous and Mesostructured Materials for Optical Applications" Chem. Mater. 2001, 13, 3140-3150. 126 Hartmann, M. "Ordered Mesoporous Materials for Bioadsorption and Biocatalysis" Chem. Mater. 2005, 17, 4577-4593. 127 Hoffmann, F., Cornelius, M., Morell, J. and Fröba, M. "Silica-Based Mesoporous Organic-Inorganic Hybrid Materials" Angew. Chem., Int. Ed. 2006, 45, 3216-3251.

133

producing silica walls templated around the surfactant micelles that resulted in ordered

mesoporous silica films.128 It was also discovered that mesoporous silica films can be

produced at the interfaces of air-water129 or water-oil systems by using surfactant

micelles as a template to grow mesoporous silica at these interfaces.130 However, it was

found that synthesis of mesoporous silica films using surfactant template had long

deposition times, the film texture was grainy131, and many times the orientation of the

channels was parallel to the substrate.132 The next advance was evaporation-induced self-

assembly (EISA).133,134,135,136 This process was based on silica/surfactant self-assembly

into thin-film mesophases during the evaporation of the solvent and of the alcohol

released by processing after inkjet printing, spin-coating, or dip-coating the substrate.

The EISA approach produced a 3D network of interconnected pores across the film.

However, the limitations on mass transport inherent to a tortuous path were not

eliminated. In addition EISA requires flat surfaces for spin- and dip-coating.

An overview of synthesis methods that produce mesoporous metal-oxide and

other oxide films with pore orientation perpendicular to the substrate surface with the use

of surfactant templating and dimensional confinement has been presented137. In

particular, a porous alumina membrane can be used to provide a dimensional

confinement for the self-assembly process resulting in nanocomposite films with pores 128 Kresge, C. T., Leonowicz, M. E., Roth, W. J., Vartuli, J. C. and Beck, J. S. "Ordered mesoporous molecular sieves synthesized by a liquid-crystal template mechanism" Nature 1992, 359, 710-712. 129 Yang, H., Coombs, N., Sokolov, I. and Ozin, G. A. "Free-standing and oriented mesoporous silica films grown at the air-water interface" Nature 1996, 381, 589-592. 130 Schacht, S., Huo, Q., Voigt-Martin, I. G., Stucky, G. D. and Schüth, F. "Oil-Water Interface Templating of Mesoporous Macroscale Structures" Science 1996, 273, 768-771. 131 Nicole, L., Boissi, egrave, re, C., eacute, dric, Grosso, D., Quach, A., Sanchez, C. and ment "Mesostructured hybrid organic & inorganic thin films" J. Mater. Chem. 2005, 15, 3598-3627. 132 Yang, H., Kuperman, A., Coombs, N., Mamiche-Afara, S. and Ozin, G. A. "Synthesis of oriented films of mesoporous silica on mica" Nature 1996, 379, 703-705. 133 Lu, Y., Ganguli, R., Drewien, C. A., Anderson, M. T., Brinker, C. J., Gong, W., Guo, Y., Soyez, H., Dunn, B., Huang, M. H. and Zink, J. I. "Continuous formation of supported cubic and hexagonal mesoporous films by sol-gel dip-coating" Nature 1997, 389, 364-368. 134 Brinker, C. J., Lu, Y., Sellinger, A. and Fan, H. "Evaporation-Induced Self-Assembly: Nanostructures Made Easy" Adv. Mat. 1999, 11, 579-585. 135 Grosso, D., Cagnol, F., Soler-Illia, G., thinsp, J, de, A, Crepaldi, E., L, Amenitsch, H., Brunet-Bruneau, A., Bourgeois, A. and Sanchez, C. "Fundamentals of Mesostructuring Through Evaporation-Induced Self-Assembly" Adv. Funct. Mater. 2004, 14, 309-322. 136 Minoofar, P. N., Dunn, B. S. and Zink, J. I. "Multiply Doped Nanostructured Silicate Sol-Gel Thin Films: Spatial Segregation of Dopants, Energy Transfer, and Distance Measurements" J. Am. Chem. Soc. 2005, 127, 2656-2665. 137 Brinker, C. J. and Dunphy, D. R. "Morphological control of surfactant-templated metal oxide films" Curr. Opin. Colloid In. 2006, 11, 126-132.

134

perpendicular to the membrane’s transverse channels.138,139 However, due to lack of a

non-nanostructured control definitive conclusions on the morphology of obtained films

cannot be made.

The concept of controlling the orientation of morphology of the mesoporous films

with the use of surfactants was theoretically predicted and subsequently demonstrated by

dip-coating silica precursors and surfactants onto modified glass slides140. Scaling of

domain size and rientation control of cylindrical domains in porous organosilicate thin

films from simple binary mixtures of an amphiphilic block copolymer and an oligomeric

organosilicate precursor (thereby forming cylindrical or spherical pores with parallel and

normal orientation to the substrate surface) has been demonstrated.141 Pore orientation

normal to the substrate surface was achieved by using Au nanoparticles to catalyze high

temperature growth of the mesoporous silica films. While the latter method provides a

means to control the pore orientation it requires a long times to form the desired films;

the presence of both parallel and normal pore orientation was reported. Another advance

into controlling pore orientation in mesoporous silica films was the use of a photo-

crosslinkable polymer film as a template. Both parallel and normal cylindrical pore

orientation was demonstrated.142 The predominance of either orientation was controlled

to 93%, and it was noted that the thinner samples (on the order of 40nm) had the best

orientation integrity. While the orientation problem was addressed in several

methods,143,144,145 some degree of pore orientation nonuniformity was retained.

138 Lu, Q., Gao, F., Komarneni, S. and Mallouk, T. E. "Ordered SBA-15 Nanorod Arrays Inside a Porous Alumina Membrane" J. Am. Chem. Soc. 2004, 126, 8650-8651. 139 Yamaguchi, A., Uejo, F., Yoda, T., Uchida, T., Tanamura, Y., Yamashita, T. and Teramae, N. "Self-assembly of a silica-surfactant nanocomposite in a porous alumina membrane" Nat. Mater. 2004, 3, 337-341. 140 Koganti, V. R. and Rankin, S. E. "Synthesis of Surfactant-Templated Silica Films with Orthogonally Aligned Hexagonal Mesophase" J. Phys. Chem. B 2005, 109, 3279-3283. 141 Freer, E. M., Krupp, L. E., Hinsberg, W. D., Rice, P. M., Hedrick, J. L., Cha, J. N., Miller, R. D. and Kim, H. C. "Oriented Mesoporous Organosilicate Thin Films" Nano Lett. 2005, 5, 2014-2018. 142 Fukumoto, H., Nagano, S., Kawatsuki, N. and Seki, T. "Photo-Alignment Behavior of Mesoporous Silica Thin Films Synthesized on a Photo-Cross-Linkable Polymer Film" Chem. Mater. 2006, 18, 1226-1234. 143 Koganti and Rankin "Synthesis of Surfactant-Templated Silica Films with Orthogonally Aligned Hexagonal Mesophase". 144 Freer, Krupp, Hinsberg, Rice, Hedrick, Cha, Miller and Kim "Oriented Mesoporous Organosilicate Thin Films". 145 Fukumoto, Nagano, Kawatsuki and Seki "Photo-Alignment Behavior of Mesoporous Silica Thin Films Synthesized on a Photo-Cross-Linkable Polymer Film".

135

Another approach to obtaining mesoporous silica films with normal pore

orientation is to combine electrochemically induced self-assembly of surfactants at the

solid-liquid interface146 and an electrochemically assisted deposition method to produce

sol-gels.147,148 This approach was demonstrated to produce highly ordered mesoporous

silica films with pore channels oriented perpendicular to the solid substrate material.149

This method is different from EISA because the mesoporous thin-film growth is guided

by both the surfactant-templated surface and sol-gel controlled by the potential of the

substrate electrode. It is also different from cathodic deposition of ordered metal or

metal-oxide films around polymeric templates150 and electrochemical precipitation of

metal oxides using self-assembly of surfactant molecules in that the electrogenerated

species in this process is a catalyst rather than a component of the final material151.

There is a recognized, pressing demand for rugged, miniaturized and portable

biosensing devices across fields as diverse as in vivo medical monitoring, disease

diagnostics, bioprocess and environmental monitoring, food and drug quality control,

genomics, and proteonomics. Thus, sol-gel films are of importance in the design of such

sensors, but they are limited by the difficulty of controlling their pore structure.

Therefore, the goal of this work is to investigate the design and utility of

electrochemically induced self-assembly of surfactants at the solid-liquid interface to

produce functionalized films with perpendicular nanocapillaries. The present study

reports the electrochemically assisted processing of organically modified silica (ormosil)

films with mesopores normal to the electrode surface (EPONs). This general fabrication

146 Choi, K. S., McFarland, E. W. and Stucky, G. D. "Electrocatalytic Properties of Thin Mesoporous Platinum Films Synthesized Utilizing Potential-Controlled Surfactant Assembly" Adv. Mat. 2003, 15, 2018-2021. 147 Shacham, R., Avnir, D. and Mandler, D. "Electrodeposition of Methylated Sol-Gel Films on Conducting Surfaces" Adv. Mat. 1999, 11, 384-388. 148 Sibottier, E., Sayen, S., eacute, phanie, Gaboriaud, F. and Walcarius, A. "Factors Affecting the Preparation and Properties of Electrodeposited Silica Thin Films Functionalized with Amine or Thiol Groups" Langmuir 2006, 22, 8366-8373. 149 Walcarius, A., Sibottier, E., Etienne, M. and Ghanbaja, J. "Electrochemically assisted self-assembly of mesoporous silica thin films" Nat. Mater. 2007, 6, 602-608. 150 Bartlett, P. N., Birkin, P. R. and Ghanem, M. A. "Electrochemical deposition of macroporous platinum, palladium and cobalt films using polystyrene latex sphere templates" Chem. Commun. 2000, 2000, 1671-1672. 151 Choi, K.-S., Lichtenegger, H. C., Stucky, G. D. and McFarl, E. W. "Electrochemical Synthesis of Nanostructured ZnO Films Utilizing Self-Assembly of Surfactant Molecules at Solid−Liquid Interfaces" J. Am. Chem. Soc. 2002, 124, 12402-12403.

136

method has been previously reported152. The solution was prehydrolyzed for 1 hour to

allow organization of the CTAB on the electrode. This organized structure has spaces

between the CTAB molecules. The electrode is placed in an unbuffered electrolyte that

contains a sol. During a brief cathodization electrolyte hydroxyl ions are produced that

catalyze the polycondensation of the sol to form a thin silica film. The sol-gel is grown

in the aqueous spaces around the CTAB hydrophilic head group in the supramolecular

assembly. Given that the CTAB structure is perpendicular to the substrate surface, the

deposited sol-gel is also perpendicular to the substrate surface. The CTAB is leached

from the film, resulting in a sol-gel coated surface in which the perpendicular

mesoporous structure is maintained. In the reported study153, tetrethylorthosilicate

(TEOS) was the sol.

In this paper, the general EPON frication is extended by organically modifying

the film precursor with a fluorescent functional group.

A2-2. Experimental

The organically modified precursor was prepared by reacting 7-hydroxy-4-methyl

coumarin (coumarin 4) with 3-(triethoxysilyl) propylisocyanate (TriMeOSiC) to produce

4-methylcoumarin-7-yl 3-(trimethoxysilyl)propylcarbamate. Analogous to a previous

report (27), partial hydrolysis of TEOS and TriMeOSiC was performed at pH 3. For

example, 13.6 mmol TEOS; 20 mL ethanol; 20 mL aqueous solution of 0.1 M NaNO3;

10-3 M HCl, and an aliquot of 10-3 M of TriMeOSiC were mixed with 4.35 mmol CTAB

(98%, Fluka) under stirring. This amount of CTAB corresponds to a CTAB/TEOS ratio

of 0.32 (higher or lower ratios give less ordered mesostructures or even non-ordered

materials.154 The mixture was reacted for 2.5 h under stirring before use. An indium tin

oxide (ITO) electrode, which served as substrate, was immersed in the above solution,

and under quiescent conditions, a cathodic potential (e.g. -1.2 to -2.1 V vs Ag/AgCl) was

applied for a defined period (typically 1–10 s, depending on the desired film thickness).

152 Walcarius, Sibottier, Etienne and Ghanbaja "Electrochemically assisted self-assembly of mesoporous silica thin films". 153 Walcarius "Electrochemically assisted self-assembly of mesoporous silica thin films". 154 Walcarius "Electrochemically assisted self-assembly of mesoporous silica thin films".

137

Under these conditions the reduction of water occurs; the OH- released catalyzes the sol-

gel reaction specifically at the electrode surface, thereby forming a functionalized silica

film. The electrode was removed from the solution immediately after the electrolysis and

rinsed with water to avoid any further deposition through an undesired evaporation-

induced condensation process. Blank experiments, performed under the same conditions

but without applying any potential to the electrode, did not yield a film on the ITO. The

samples were dried at room temperature overnight. In some cases, they were dried

further at 110°C. The reported drying at 130°C 155 was not used to avoid possible

thermal decomposition of 4-methylcoumarin-7-yl 3’-trimethoxysilyl)propylcarbamate.

The dried samples were extracted overnight with 10 mL of 50:50 ethanol 0.1M HCl (aq.)

to extract CTAB as well as residual sol.

All fluorescence measurements were obtained with a Horiba Jobin Yvon

Fluorolog-3-22 spectrofluorometer. AFM images were obtained with a Agilent 5500

AFM Microscope.

A2-3. Results and Discussion

An initial screening of the fluorescence from films on ITO is shown in Figure A2-

1. Under UV radiation (Fig. A2-1b), fluorescence is observed. In addition, the EPON-

coated ITO appeared white due to light scattering. Note that the sample used in Figure

A2-1 employed an extended deposition time so that a layer of silica spheres was

deposited on top of the EPON. This procedure was used to achieve visual evidence of

incorporation of coumarin 4 into the sol-gel structures.

Figure A2-1. (a) Left sample—Blank ITO Electrode, right sample—ITO Electrode with the EPON film. (b) Samples in same position under UV light, fluorescence is observed for ITO Electrode with EPON film

155 Walcarius "Electrochemically assisted self-assembly of mesoporous silica thin films".

(a) (b)

138

AFM images of the EPON on ITO are shown in Figure A2-2. Here the

electrochemically assisted deposition was 10 s at -2.1 V vs Ag/AgCl. First, it is apparent

that the coverage of the ITO is not uniform; instead, the EPON forms as islands on the

surface. This pattern is not surprising in that when used as received ITO surfaces exhibit

non-uniform catalysis, so the electrolysis of water, and therefore the pH, will vary across

the electrode surface, as shown by Figure A2-2. Second, within a given island, the film

thickness is uniform, with the deposted mesopores normal to the ITO electrode surface as

can be seen in Figure A2-3. The EPON pore structure previously reported with a TEOS

precursor156 clearly is retained with this fluorescing film prepared from TEOS,

TriMeOSiC mixed sol. Also demonstrated was that, consistent with a TEOS precursor,

the thickness of the film depended upon electroysis time and the applied potential.

Figure A2-2. EPON-Coated ITO Electrode, plating time 10 s at -2.1V. Imaging of the plating interface shows the difference in surface morphology that of the ITO and that of EPON film, which indicates EPON film is deposited on the ITO surface

156 Walcarius "Electrochemically assisted self-assembly of mesoporous silica thin films".

139

Figure A2-3. EPON-coated ITO electrode, plating time: (a) 30 s -2.1V. The deposited film indicates normal to the electrode surface orientation of the deposited EPON film. (b) EPON-coated ITO electrode, magnification of (a) reveals normal orientation of the mesopores

The UV-Vis and fluorescence spectra of 7-hydroxy-4-methyl coumarin and 4-

methylcoumarin-7-yl 3’-(trimethoxysilyl)propylcarbamate solutions of the same

concentration (10 μmol/L) produced identical intensity emission at 387 nm. In addition

the absorbance and excitation spectra were the same, the absorbance maximum was at

322 nm in both cases. This suggests that the incorporation of the silyl group did not

affect the electronic excitation and emission mechanisms, an observation consistent with

prior reports on siloxy-derivatized coumarin 4 that was incorporated into sol-gel

monoliths.157,158, 159

The incorporation of coumarin 4 in the sol-gel film is stable after the initial

washing to remove the unbound fluorophore. For example, Figure A2-4 compares the

fluorescence from a film at various stages of fabrication. After washing the coumarin-4

emission drops by a factor of 2.2, indicating the presence of non-bonded 4-

methylcoumarin-7-yl 3- (trimethoxysilyl) - propylcarbamate. Further washing did not

change the spectra. Wetting the sample with water as well as organic solvents such as n-

157 Suratwala, T., Gardlund, Z., Davidson, K., Uhlmann, D. R., Watson, J. and Peyghambarian, N. "Silylated Coumarin Dyes in Sol-Gel Hosts. 1. Structure and Environmental Factors on Fluorescent Properties" Chem. Mater. 1998, 10, 190-198. 158 Suratwala, T., Gardlund, Z., Davidson, K., Uhlmann, D. R., Watson, J., Bonilla, S. and Peyghambarian, N. "Silylated Coumarin Dyes in Sol-Gel Hosts. 2. Photostability and Sol-Gel Processing" Chem. Mater. 1998, 10, 199-209. 159 Suratwala, T., Gardlund, Z., Boulton, J. M. and Uhlmann, D. R. Series Incorporation of triethoxysilyl functionalized Coumarin 4 in sol-gel hosts, 13 October, 1994 Conference.

(a) (b)

140

hexane makes the opaque, light-scattering sample translucent with an accompanying

fluorescence emission decrease by a factor of 7.3. After the water or organic solvent

evaporates the emission returns to the original intensity obtained after washing procedure.

This shows that the synthesized 4-methylcoumarin-7-yl 3’-(trimethoxysilyl)propyl-

carbamate is permanently bound into the EPON material.

Figure A2-4. (a) Dry Sol-Gel Film; (b) Dry, after CTAB and free polymer extraction; (c) Wetted with water; (d) Dry. (b) and (d) overlap. The difference in emission intensity between (a) and (b) amounts to removed 4-methylcoumarin-7-yl 3-(trimethoxysilyl) propylcarbamate not bound to EPON film. (b) and (d) are the same film before and after wetting, where (c) shows intensity drop when the obtained film is washed of CTAB and 4-methylcoumarin-7-yl 3-(trimethoxysilyl)propylcarbamate, that is not bound; dried and wetted again

The excitation and emission spectra of the TriMeOSiC-containing EPON were

obtained at pH 1, 2.2, and 13.3 as well as in distilled water (Figures A2-5 and A2-6). The

shifts of the fluorescence spectra with pH are consistent with those for coumarin 4 in

homogeneous solutions. This observation demonstrates that the film on ITO is

141

permeable and that coumarin 4 sites are on the phase boundary between the pore volume

and the bulk film. These are requirements for applying this fabrication method to

development of an optical sensor.

Figure A2-5. Overlaid excitation spectra of EPON film subjected to different pH: (a) DI water; (b) pH 1; (c) pH 2.2; (d) pH 13.3. Excitation maximum and excitation peak shape shifts based on pH of the wetting solution

Finally, the results in Figure A2-5 and A2-6 were not changed when drying

overnight at room temperature. This is in contrast to comparable experiments performed

on monolilthic sol-gels.160,161,162,163 With monolithic, the aging process, which is

160 Suratwala, Gardlund, Davidson, Uhlmann, Watson and Peyghambarian "Silylated Coumarin Dyes in Sol-Gel Hosts. 1. Structure and Environmental Factors on Fluorescent Properties". 161 Suratwala, Gardlund, Davidson, Uhlmann, Watson, Bonilla and Peyghambarian "Silylated Coumarin Dyes in Sol-Gel Hosts. 2. Photostability and Sol-Gel Processing". 162 Oh, E. O., Gupta, R. K. and Whang, C. M. "Effects of pH and Dye Concentration on the Optical and Structural Properties of Coumarin-4 Dye-Doped SiO2-PDMS Xerogels" J. Sol-Gel Sci. Technol. 2003, 28, 279-288.

142

accompanied by drying, causes a rigidochromic shift in the fluorescence spectra of

dopants. With a nanoscale film, this aging/drying step is accelerated. From the

agreement of the spectra, apparently sol-gel processing is virtually complete in ca. 24 h,

whether it is performed at room temperature or in an oven.

Figure A2-6. Overlaid emission spectra of EPON film subjected to different pH: (a) DI water; (b) pH 1; (c) pH 2.2; (d) pH 13.3 Emission maximum and peak shape, consistent with excitation peak changes, shifts based on pH of the wetting solution

A2-4. Conclusions

The electrochemically assisted sol-gel processing method that is reported herein

yields a fluorescent, nanoscale film. The obtained EPON films have low tortuosity and

mesopores that are normal to the electrode surface. The synthesized 4-methylcoumarin-

7-yl 3-(trimethoxysilyl)propyl-carbamate sites are stabilized in the EPON by covalent

163 Oh, E. O., Gupta, R. K., Cho, N.-H., Yoo, Y.-C., Cho, W.-S. and Whang, C. M. "Influence of pH and Dye Concentration on the Physical Properties and Microstructure of New Coumarin 4 Doped SiO2-PDMS ORMOSIL " Bull. Korean Chem. Soc. 2003, 24, 299-306.

143

bonding and are sensitive to pH. The emission maximum shifts based on changes in pH

and are consistent with the spectra for coumarin 4 in homogeneous solution. A potential

limitation which we are studying further is that the intensity decreases when the EPON is

wetted. Nanometer scale films are expected to have fast response times especially with

the observed EPON structure, which in combination with the sensing fluorophore

provides a basis for practical optical sensor platforms.

A2.5 Acknowledgements

This work was in-part sponsored by The Ohio Third Frontier IDCAST Wright

Center for Innovation.

144

References:

Adam, L.; Gordon, G.; Pierce, D. Bleach 2001 Predictive Model, Miami University, Oxford, Ohio. Copyright (c) 2001 AwwaRF and AWWA.

Adam, L. C. An Investigation of Factors Involved In The Decomposition of Sodium

Hypochlorite Ph.D. Dissertation, Miami University, Oxford, OH, 1994. Adam, L. C.; Fabian, I.; Suzuki, K.; Gordon, G. "Hypochlorous Acid Decomposition in

the pH 5-8 Region" Inorg. Chem. 1992, 31, 3534-3541. Adam, L. C.; Gordon, G. "Direct and Sequential Potentiometric Determination of

Hypochlorite, Chlorite, and Chlorate Ions When Hypochlorite Ion is Present in Large Excess" Anal. Chem. 1995, 67, 535-540.

Adam, L. C.; Gordon, G. "Hypochlorite ion decomposition: Effects of temperature, ionic

strength, and chloride ion" Inorg. Chem. 1999, 38, 1299-1304. Asami, M.; Kosaka, K.; Kunikane, S. "Bromate, chlorate, chlorite and perchlorate in

sodium hypochlorite solution used in water supply" J. Water Supply Res. Technol. 2009, 58, 107-115.

Ayers, G. H. B., M. H. "Catalytic Decomposition of Hypochlorite Solution by Iridium

Compounds. II. Kinetic Studies" J. Am. Chem. Soc. 1955, 77, 828-833. Bartlett, P. N.; Birkin, P. R.; Ghanem, M. A. "Electrochemical deposition of macroporous

platinum, palladium and cobalt films using polystyrene latex sphere templates" Chem. Commun. 2000, 2000, 1671-1672.

Bouland, S.; Duguet, J. P.; Montiel, A. "Evaluation of bromate ions level introduced by

sodium hypochlorite during postdisinfection of drinking water" Environ. Technol. 2005, 26, 121-125.

Brinker, C. J.; Dunphy, D. R. "Morphological control of surfactant-templated metal oxide

films" Curr. Opin. Colloid In. 2006, 11, 126-132. Brinker, C. J.; Lu, Y.; Sellinger, A.; Fan, H. "Evaporation-Induced Self-Assembly:

Nanostructures Made Easy" Adv. Mat. 1999, 11, 579-585. CDHS, "Determination of Perchlorate by Ion Chromatography", California Department

of Health Services (CDHS), 1997, Retrieved from: http://www.cdph.ca.gov/certlic/drinkingwater/Documents/Perchlorate/SRLperchloratemethod1997.pdf

145

CDPH, "History of Perchlorate in California Drinking Water", California Department of Public Health (CDPH), 2007, Retrieved 07/28/09, from: http://www.cdph.ca.gov/certlic/drinkingwater/Pages/Perchloratehistory.aspx

Chlorine Institute, "Bromate in Sodium Hypochlorite Potable Water Treatment", 2004,

Retrieved 04/14/09, from: http://www.chlorineinstitute.org/files/FileDownloads/BromateinNaOCl-PotableWaterTreatment.pdf

Chlorine Institute, "Pamthlet 96 Sodium Hypochlorite Manual, Edition 3," Chlorine

Institute, 2006, Retrieved 07/08/09, from: http://www.chlorineinstitute.org/files/Pamphlet096Edition3April2006FinalWebsite.pdf

Choi, K.-S.; Lichtenegger, H. C.; Stucky, G. D.; McFarl, E. W. "Electrochemical

Synthesis of Nanostructured ZnO Films Utilizing Self-Assembly of Surfactant Molecules at Solid-Liquid Interfaces" J. Am. Chem. Soc. 2002, 124, 12402-12403.

Choi, K. S.; McFarland, E. W.; Stucky, G. D. "Electrocatalytic Properties of Thin

Mesoporous Platinum Films Synthesized Utilizing Potential-Controlled Surfactant Assembly" Adv. Mat. 2003, 15, 2018-2021.

Choi, Y.-K.; King, T.-J.; Hu, C. "Nanoscale CMOS spacer FinFET for the terabit era"

Electron. Devic. Lett. 2002, 23, 25-27. Clesceri, L. S. G., A. E.; Eaton, A. D. Eds. Standard Methods for the Examination of

Water and Wastewater 20th ed.; American Public Health Association: Washington, DC, 1998.

Creed, J. T. B., C.A.; Martin, T.D. U.S. Environmental Agency Method 200.8 Revison

5.4 1994. Davies, C. W. "397. The extent of dissociation of salts in water. Part VIII. An equation

for the mean ionic activity coefficient of an electrolyte in water, and a revision of the dissociation constants of some sulphates" J. Chem. Soc. 1938, 1938, 2093-2098.

Davis, C. O. Perchlorate explosives, E. I. du Pont de Nemours & Co., 1940, USA US

2190703 Espenson, J. H. Chemical Kinetics and Reaction Mechanisms, 2nd ed.; McGraw-Hill,

Inc., 1995. Farkas, L.; Lewin, M.; Bloch, R. "The Reaction between Hypochlorite and Bromides" J.

Am. Chem. Soc. 1949, 71, 1988-1991.

146

Freer, E. M.; Krupp, L. E.; Hinsberg, W. D.; Rice, P. M.; Hedrick, J. L.; Cha, J. N.; Miller, R. D.; Kim, H. C. "Oriented Mesoporous Organosilicate Thin Films" Nano Lett. 2005, 5, 2014-2018.

Fukumoto, H.; Nagano, S.; Kawatsuki, N.; Seki, T. "Photo-Alignment Behavior of

Mesoporous Silica Thin Films Synthesized on a Photo-Cross-Linkable Polymer Film" Chem. Mater. 2006, 18, 1226-1234.

Gordon, G.; Adam, L. C.; Bubnis, B.; Hoyt, B.; Gillette, S. J.; Wilczak, A. "Controlling

the formation of chlorate ion in liquid hypochlorite feedstocks" J. Am. Water Work Assoc. 1993, 85, 89-97.

Gordon, G.; Cooper, W. J.; Rice, R. G.; Pacey, G. E. Disinfectant Residual Measurement

Methods, Second ed.; AWWA Research Foundation and American Water Works Association: Denver, 1992.

Greer, M. A., Goodman G. G., Pleus R. C, Greer S. E. "Health effects assessment for

environmental perchlorate contamination: the dose response for inhibition of thyroidal radioiodine uptake in humans" Environ. Health Perspect. 2002, 110, 927-937.

Greiner, P.; Mclellan, C.; Bennet, D.; Ewing, A. "Occurrence of perchlorate in sodium

hypochlorite" J. Am. Water Work Assoc. 2008, 100, 68-74. Grosso, D.; Cagnol, F.; Soler-Illia, G.; thinsp; J; de; A; Crepaldi, E.; L; Amenitsch, H.;

Brunet-Bruneau, A.; Bourgeois, A.; Sanchez, C. "Fundamentals of Mesostructuring Through Evaporation-Induced Self-Assembly" Adv. Funct. Mater. 2004, 14, 309-322.

Harned, H. S.; Owen, B. B. The Physical Chemistry of Electrolytic Solutions, 2nd ed.;

Reinhold Publishing Corporation: New York, 1950. Harris, D. C. Quantitative Chemical Analysis, 5th ed.; W.H. Freeman and Company:

New York, 2001. Hartmann, M. "Ordered Mesoporous Materials for Bioadsorption and Biocatalysis"

Chem. Mater. 2005, 17, 4577-4593. Hedrick, E. B., T.;Slingsby, R.;Munch, D. U.S. Environmental Agency Method 332.0

Revison 1.0 2005. Hoffmann, F.; Cornelius, M.; Morell, J.; Fröba, M. "Silica-Based Mesoporous Organic-

Inorganic Hybrid Materials" Angew. Chem., Int. Ed. 2006, 45, 3216-3251. Housecroft, C. E.; Sharpe, A. G. Inorganic Chemistry; Pearson Education Limited, 2001.

147

Ikeda, Y.; Tang, T.; Gordon, G. "Iodometric method for determination of trace chlorate ion" Anal. Chem. 1984, 56, 71-73.

Jackson, P. E.; Gokhale, S.; Streib, T.; Rohrer, J. S.; Pohl, C. A. "Improved method for

the determination of trace perchlorate in ground and drinking waters by ion chromatography" J. Chromatogr., A 2000, 888, 151-158.

Kang, N.; Anderson, T. A.; Andrew Jackson, W. "Photochemical formation of

perchlorate from aqueous oxychlorine anions" Anal. Chim. Act. 2006, 567, 48-56. Kang, N.; Jackson, W. A.; Dasgupta, P. K.; Anderson, T. A. "Perchlorate production by

ozone oxidation of chloride in aqueous and dry systems" Sci. Total Environ. 2008, 405, 301-309.

Keith, J. D.; Pacey, G. E.; Cotruvo, J. A.; Gordon, G. "Experimental results from the

reaction of bromate ion with synthetic and real gastric juices" Toxicology 2006, 221, 225-228.

Koganti, V. R.; Rankin, S. E. "Synthesis of Surfactant-Templated Silica Films with

Orthogonally Aligned Hexagonal Mesophase" J. Phys. Chem. B 2005, 109, 3279-3283.

Kresge, C. T.; Leonowicz, M. E.; Roth, W. J.; Vartuli, J. C.; Beck, J. S. "Ordered

mesoporous molecular sieves synthesized by a liquid-crystal template mechanism" Nature 1992, 359, 710-712.

Krozer, A.; Rodahl, M. "X-ray photoemission spectroscopy study of UV/ozone oxidation

of Au under ultrahigh vacuum conditions" J. Vac. Sc. Technol. A. 1997, 15, 1704-1709.

Li, Y. E., J. George "Analysis of Perchlorate in Water by Reversed-Phase LC/ESI-

MS/MS Using an Internal Standard Technique" Anal. Chem. 2005, 77, 4453-4458.

Liebhafsky, H. A.; Mohammad, A. "A Third-order Ionic Reaction without Appreciable

Salt Effect" J. Phys. Chem. 1934, 38, 857-866. Lister, M. W. "Decomposition of Sodium Hypochlorite: The Catalyzed Reaction" Can. J.

Chem. 1956, 34, 479–488. Lister, M. W. "Decomposition of Sodium Hypochlorite: The Uncatalyzed Reaction" Can.

J. Chem. 1956, 34, 465-478. Liu, W., Andrews, S. A., Stefan, M. I., & Bolton, J. R. "Optimal methods for quenching

H2O2 residuals prior to UFC testing" Water Res. 2003, 37, 3697-3703.

148

Lu, Q.; Gao, F.; Komarneni, S.; Mallouk, T. E. "Ordered SBA-15 Nanorod Arrays Inside a Porous Alumina Membrane" J. Am. Chem. Soc. 2004, 126, 8650-8651.

Lu, Y.; Ganguli, R.; Drewien, C. A.; Anderson, M. T.; Brinker, C. J.; Gong, W.; Guo, Y.;

Soyez, H.; Dunn, B.; Huang, M. H.; Zink, J. I. "Continuous formation of supported cubic and hexagonal mesoporous films by sol-gel dip-coating" Nature 1997, 389, 364-368.

Mack, E. L. "Electrolytic Formation of Perchlorate" J. Phys. Chem. 1917, 21, 238-264. MassDEP, "Perchlorate in Public Drinking Water", 2006, Retrieved 07/21/09, from:

http://www.mass.gov/dep/toxics/pchlorqa.htm Metz, S.; Trautmann, C.; Bertsch, A.; Renaud, P. "Polyimide microfluidic devices with

integrated nanoporous filtration areas manufactured by micromachining and ion track technology" J. Micromech. Microeng. 2004, 14, 324-331.

Miller, K. G.; Pacey, G. E.; Gordon, G. "Automated iodometric method for determination

of trace chlorate ion using flow injection analysis" Anal. Chem. 1985, 57, 734-737.

Minoofar, P. N.; Dunn, B. S.; Zink, J. I. "Multiply Doped Nanostructured Silicate Sol-

Gel Thin Films: Spatial Segregation of Dopants, Energy Transfer, and Distance Measurements" J. Am. Chem. Soc. 2005, 127, 2656-2665.

Motzer, W. "Perchlorate: Problems, Detection, and Solutions" Environ. Forensics 2001,

2, 301-311. Nicole, L.; Boissi; egrave; re, C.; eacute; dric; Grosso, D.; Quach, A.; Sanchez, C.; ment

"Mesostructured hybrid organic & inorganic thin films" J. Mater. Chem. 2005, 15, 3598-3627.

Oh, E. O.; Gupta, R. K.; Cho, N.-H.; Yoo, Y.-C.; Cho, W.-S.; Whang, C. M. "Influence

of pH and Dye Concentration on the Physical Properties and Microstructure of New Coumarin 4 Doped SiO2-PDMS ORMOSIL " Bull. Korean Chem. Soc. 2003, 24, 299-306.

Oh, E. O.; Gupta, R. K.; Whang, C. M. "Effects of pH and Dye Concentration on the

Optical and Structural Properties of Coumarin-4 Dye-Doped SiO2-PDMS Xerogels" J. Sol-Gel Sci. Technol. 2003, 28, 279-288.

Perlmutter-Hayman, B.; Stein, G. "The Kinetics of the Decomposition of Alkaline

Solutions of Hypobromite-Specific Ionic Effects on Reaction Rate" J. Phys. Chem. 1959, 63, 734-738.

149

Puckett, S. D.; Heuser, J. A.; Keith, J. D.; Spendel, W. U.; Pacey, G. E. "Interaction of ozone with gold nanoparticles" Talanta 2005, 66, 1242-1246.

Renner, R. "Study finding perchlorate in fertilizer rattles industry" Environ. Sci. Technol

1999, 33, 394A-395A. Renner, R. "EPA perchlorate decision flawed, say advisers" Environ. Sci. Technol. 2009,

43, 553-554. Robinson, R. A.; Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths Publications

Limited: London, 1959. Roehl, R.; Slingsby, R.; Avdalovic, N.; Jackson, P. E. "Applications of ion

chromatography with electrospray mass spectrometric detection to the determination of environmental contaminants in water" J. Chromatogr., A 2002, 956, 245-254.

Routt, J.; Mackey, E.; Noack, R. "Committee Report: Disinfection Survey, Part 2 -

Alternatives, experiences, and future plans" J. Am. Water Work Assoc. 2008, 100, 110-124.

Saliba, N.; Parker, D. H.; Koel, B. E. "Adsorption of oxygen on Au(111) by exposure to

ozone" Surf. Sci. 1998, 410, 270-282. Salov, V. V. Y., J.; Shibata, Y.; Morita, M. "Determination of inorganic Halogen Species

by Liquid Chromatography with Inductively Coupled Argon Plasma Mass Spectrometry" Anal. Chem. 1992, 64, 2425-2428.

Sanchez, C.; Julian, B.; Belleville, P.; Popall, M. J. Mater. Chem. 2005, 15, 3559. Sato, H., Homma, T. "Fabrication of high-aspect-ratio arrayed structures using Si

electrochemical etching" Sci. Technol. Adv. Mat. 2006, 7, 468-474. Schacht, S.; Huo, Q.; Voigt-Martin, I. G.; Stucky, G. D.; Schüth, F. "Oil-Water Interface

Templating of Mesoporous Macroscale Structures" Science 1996, 273, 768-771. Scott, B. J.; Wirnsberger, G.; Stucky, G. D. "Mesoporous and Mesostructured Materials

for Optical Applications" Chem. Mater. 2001, 13, 3140-3150. Shacham, R.; Avnir, D.; Mandler, D. "Electrodeposition of Methylated Sol-Gel Films on

Conducting Surfaces" Adv. Mat. 1999, 11, 384-388. Shah, J.; Qureshi, N. "Chlorine Gas vs. Sodium Hypochlorite: What's the Best Option"

Opflow 2008, 24-27.

150

Sibottier, E.; Sayen, S.; eacute; phanie; Gaboriaud, F.; Walcarius, A. "Factors Affecting the Preparation and Properties of Electrodeposited Silica Thin Films Functionalized with Amine or Thiol Groups" Langmuir 2006, 22, 8366-8373.

Snyder, S. A.; Pleus, R. C.; Vanderford, B. J.; Holady, J. C. "Perchlorate and chlorate in

dietary supplements and flavor enhancing ingrediants" Anal. Chim. Act. 2006, 567, 23-32.

Snyder, S. A.; Vanderford, B. J.; Rexing, D. J. "Trace analysis of bromate, chlorate,

iodate, and perchlorate in natural and bottled waters" Environ. Sci. Technol. 2005, 39, 4586-4593.

Stetson, S. J.; Wanty, R. B.; Helsel, D. R.; Kalkhoff, S. J.; Macalady, D. L. "Stability of

low levels of perchlorate in drinking water and natural water samples" Anal. Chim. Act. 2006, 567, 108-113.

Strawson, J.; Zhao, Q.; Dourson, M. "Reference dose for perchlorate based on thyroid

hormone change in pregnant women as the critical effect" Regul. Toxicol. Pharmacol. 2004, 39, 44-65.

Su, Y. S.; Morrison, D. T.; Ogle, R. A. "Chemical kinetics of calcium hypochlorite

decomposition in aqueous solutions" J. Chem. Health Safety 2009, 16, 21-25. Suratwala, T.; Gardlund, Z.; Boulton, J. M.; Uhlmann, D. R. In Sol-Gel Optics III;

Mackenzie, J. D., Ed.; SPIE, 1994; Vol. 2288, pp 310-320. Suratwala, T.; Gardlund, Z.; Davidson, K.; Uhlmann, D. R.; Watson, J.; Bonilla, S.;

Peyghambarian, N. "Silylated Coumarin Dyes in Sol-Gel Hosts. 2. Photostability and Sol−Gel Processing" Chem. Mater. 1998, 10, 199-209.

Suratwala, T.; Gardlund, Z.; Davidson, K.; Uhlmann, D. R.; Watson, J.; Peyghambarian,

N. "Silylated Coumarin Dyes in Sol-Gel Hosts. 1. Structure and Environmental Factors on Fluorescent Properties" Chem. Mater. 1998, 10, 190-198.

Sweetin, D. L. Developments in the Analytical Methodology For the Determination of

Free available Chlorine, Inorganic Chloramines, and Oxyhalogen Species Ph.D. Dissertation, Miami University, Oxford, OH, 1993.

Tikkanen, M. W. "Development of a drinking water regulation for perchlorate in

California" Anal. Chim. Act. 2006, 567, 20-25. Tomiyasu, H.; Fukutomi, H.; Gordon, G. "Kinetics and mechanism of ozone

decomposition in basic aqueous solution" Inorg. Chem. 1985, 24, 2962-2966. Urbansky, E. T. "Perchlorate as an environmental contaminant" Env. Sci. Pollut. Res.

2002, 9, 187-192.

151

Urbansky, E. T.; Brown, S. K.; Magnuson, M. L.; Kelty, C. A. "Perchlorate levels in

samples of sodium nitrate fertilizer derived from Chilean caliche" Environ. Pollut. 2001, 112, 299-302.

Urbansky, E. T.; Collette, T. W.; Robarge, W. P.; Hall, W. L.; Skillen, J. M.; Kane, P. F.

Survey of fertilizers and related materials for perchlorate (Cl04-), EPA, 2001.

Urbansky, E. T.; Magnuson, M. L.; Freeman, D.; Jelks, C. "Quantitation of perchlorate

ion by electrospray ionization mass spectrometry (ESI-MS) using stable association complexes wih organic cations and bases to enhance selectivity" J. Anal. At. Spectrom. 1999, 14, 1861-1866.

USEPA, "Known Perchlorate Releases in the U.S. - March 25th, 2005", United States

Environmental Protection Agency, 2005, Retrieved 08/01/09, from: http://www.epa.gov/swerffrr/pdf/detect0305.pdf

USEPA Draft Regulatory Determinations Support Document for Selected Contaminants

from the Second Drinking Water Contaminant Candidate List (CCL2) EPA Report 815-D-06-007, United States Environmental Protection Agency (USEPA): Washington, DC. , 2006.

USEPA, "Interim Drinking Water Health Advisory for Perchlorate", United States

Environmental Protection Agency, 2008, Retrieved 07/27/09, from: http://www.epa.gov/OGWDW/contaminants/unregulated/pdfs/healthadvisory_perchlorate_interim.pdf

USFDA, "2004-2005 Exploratory Survey Data on Perchlorate in Food", United States

Food & Drug Administration, 2005, Retrieved 08/02/09, from: http://www.fda.gov/Food/FoodSafety/FoodContaminantsAdulteration/ChemicalContaminants/Perchlorate/ucm077685.htm

USGAO Perchlorate: A System to Track Sampling and Cleanup Results Is Needed

United States Government Accountability Office (USGAO): Washington, D.C., 2005.

Walcarius, A.; Sibottier, E.; Etienne, M.; Ghanbaja, J. "Electrochemically assisted self-

assembly of mesoporous silica thin films" Nat. Mater. 2007, 6, 602-608. Weinberg, H. S.; Delcomyn, C. A.; Unnam, V. "Bromate in Chlorinated Drinking

Waters: Occurrence and Implications for Future Regulation" Environ. Sci. Technol. 2003, 37, 3104-3110.

Wendelken S. C., V. L. E., Coleman D. E., Munch D. J. U.S. Environmental Protection

Agency Method 331.0 Revision 1.0 2005.

152

WHO Guidelines for Drinking Water Quality; World Health Organization: Singapore, 2006, Retrieved from: http://www.who.int/water_sanitation_health/dwq/gdwq3rd_add1.pdf

Wood, D. W. I. Determination of Disinfectant Residuals in Chlorine Dioxide Treated

Water Using Flow Injection Analysis Ph.D. Dissertation, Miami University, Oxford, OH, 1990.

Xu, C. X.; Zhang, X. S.; Sun, X. W. "Preparation of Porous Alumina by Anodization" J.

Metastab. Nanocryst. 2005, 23, 75-78. Yamaguchi, A.; Uejo, F.; Yoda, T.; Uchida, T.; Tanamura, Y.; Yamashita, T.; Teramae,

N. "Self-assembly of a silica-surfactant nanocomposite in a porous alumina membrane" Nat. Mater. 2004, 3, 337-341.

Yang, H.; Coombs, N.; Sokolov, I.; Ozin, G. A. "Free-standing and oriented mesoporous

silica films grown at the air-water interface" Nature 1996, 381, 589-592. Yang, H.; Kuperman, A.; Coombs, N.; Mamiche-Afara, S.; Ozin, G. A. "Synthesis of

oriented films of mesoporous silica on mica" Nature 1996, 379, 703-705. Zakett, D.; Flynn, R. G. A.; Cooks, R. G. "Chlorine isotope effects in mass spectrometry

by multiple reaction monitoring" J. Phys. Chem. 1978, 82, 2359-2362.