MIAMI UNIVERSITY The Graduate School. Certificate for Approving the Dissertation

Transcription

1 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

2 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 LCMS/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 secondorder rate constant. Chapter 5 focuses on validation and application of the developed predictive expression on various hypochlorite ion solutions. Bleach 21 Predictive Model was used to predict the decomposition of hypochlorite ion, and the output was used together

3 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.

4 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 29 Dissertation Directors: Dr. Gilbert E. Pacey and Dr. Gilbert Gordon

5 Aleksey N. Pisarenko 29

6 Table of Contents List of Tables List of Figures Dedication Acknowledgements vii x xviii xix 1. Introduction Perchlorate Ion: Introduction Perchlorate Ion: Toxicity and Regulation Hypochlorite Ion Solutions as Potential Source of Perchlorate Ion Research Objectives 5 2. Analysis and Sample Preparation of Hypochlorite Ion Solutions: Analytical Methods Summary Introduction to the Analysis of Sodium Hypochlorite Solutions Transition metal ions: Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni Specific Conductance, Ionic Strength, and ph Measurements Results and discussion The LCMS/MS Analysis of Perchlorate, Bromate, and Chlorate Ions Validation of LCMS/MS Method for the Analysis of Hypochlorite Solutions Iodometric Titrations: Analysis of Chlorite, Chlorate, and Hypochlorite Ions AdamGordon Method Method Selection for the Measurement of Chlorate Ion Selection of Quenching Agent Safety, Ease of Handling, Transport, and Stability Ability to Quench Hypochlorite Ion Reproducibly Impact on the Analysis of Bromate, Chlorate, and iii

7 Perchlorate Ion Quenching Agent Selection Summary Conclusions Experimental Design: Identifying Factors Impacting the Perchlorate Ion Formation in Hypochlorite Ion Solutions Experimental Matrix and Chemicals Effect of Hypochlorite Ion Concentration Effect of Chlorate Ion Concentration Effect of Transition Metal Ions (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) Effect of Noble Metal Ions (Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ) Effect of Chlorite Ion Concentration Combined Effect of Transition Metal Ions, Chlorite and Bromide Ions Effect of Bromide Ion and Bromate Ion Concentration Effect of Ionic Strength Effect of ph Conclusions 6 4. Kinetics of Perchlorate Ion Formation and Determination of the Rate Law Reaction Order with Respect to Chlorate Ion: ln (d[clo 4 ]/dt) vs. ln [ClO 3 ] Reaction Order with Respect to Chlorate Ion: ln (d[clo 4 ]/dt) vs. ln [OCl ] Multiple Reaction Pathways Parallel Reaction Pathway Consecutive Reaction Pathway Ionic Strength Effect on the Rate of Perchlorate Ion Formation Dependence of the SecondOrder Rate Constant on the Ionic Strength 84 iv

8 4.4.2 Dependence of the SecondOrder Rate Constant on the Temperature Combining the Effects of the Ionic Strength and Temperature on the SecondOrder Rate Constant Conclusions The Perchlorate Ion Formation Model: Validation and Applications Predicted Perchlorate Ion Formation in Bulk Sodium Hypochlorite Solutions Predicted Perchlorate Ion Formation in RealWorld Bulk Sodium Hypochlorite Solutions Using Perchlorate Ion Formation Model to Determine Implications of Bulk Sodium Hypochlorite Solutions Storage Application of the Perchlorate Model to OSG Sodium Hypochlorite Solutions Application of The Perchlorate Model to Calcium Hypochlorite Solutions Potential Contribution of Perchlorate Ion to Drinking Water 114 from Various Hypochlorite Ion Solutions Conclusions Conclusions Summary Recommendation to Water Utilities 12 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 v

9 Appendix 2. Electrochemically Assisted Processing of Organically Modified, Perpendicularly Oriented Mesoporous Silica Films with Fluorescent Functionality 132 A2.1. Introduction 132 A22. Experimental Details 136 A23. Results and Discussion 137 A24. Conclusions 142 A2.5 Acknowledgements 143 References 144 vi

10 List of Tables Table 1. ICPMS 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.19 M IO 3 2 Table 5. Comparison of measurements by the LCMS/MS and iodometric titration for bulk hypochlorite ion solutions (n = 7) 21 Table 6. Comparison of measurements by the LCMS/MS and iodometric titration for OSG sodium hypochlorite solutions at less than 1. g/l ClO 3 (< 1 mm) (n 3) 22 Table 7. Effects of malonic acid (MA) on recoveries of chlorate, perchlorate, and bromate ions measured by LCMS/MS (n = 3, replicate Samples Analyzed in triplicate; S.D. = standard deviation) 3 Table 8. Effects of malonic acid (MA) on analysis of perchlorate and bromate ions at different dilutions (n = 3; S.D. = standard deviation) 3 Table 9. Effects of quenching agent on analysis of chlorate comparison of LCMS/MS and titration results (n = 3, replicate samples analyzed in triplicate; S.D. = standard deviation) 31 Table 1. Perchlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period 31 Table 11. Bromate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period 32 Table 12. Chlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period 32 Table 13. Summary of quenching agent test results and decisionmaking matrix 33 Table 14. Changes in perchlorate ion concentration of samples spiked with Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ (Noble Me) vs. control (no spike), incubated at 5 ºC 46 Table 15. Decomposition of hypochlorite ion at 3 ºC in solutions, at various initial concentrations of chlorate ion ([ClO 3 ] ) at ph ~ Table 16. Reaction order with respect to chlorate ion and corresponding vii

11 correlation coefficients in solutions at constant hypochlorite ion at ph ~ 12.5 and various temperatures 69 Table 17. Decomposition of hypochlorite ion at 3 ºC in solutions at constant chlorate ion concentration at ph ~ Table 18. Parallel reaction pathway experimental rate constants in solutions, at various hypochlorite ion and constant chlorate ion at ph ~ Table 19. Consecutive reaction pathway experimental rate constants at various initial concentrations of hypochlorite ion and constant chlorate ion at ph ~ Table 2. Ionic strength (μ) of hypochlorite ion solutions at various chlorate ion at 4 ºC experiments (TDS = Total Dissolved Solids) 81 Table 21. Ionic strength (μ) of hypochlorite ion solutions at various hypochlorite ion at 4 ºC experiments (TDS=Total Dissolved Solids) 81 Table 22. Slopes and intercepts of leastsquares lines shown in Figure Table 23. Experimental and predicted secondorder rate constants at variable ionic strength and temperature (k exp = experimental k 2 ; k pred = predicted k 2 ) 91 Table 24. Predicted changes in hypochlorite ion, chlorate ion, and d[clo 4 ]/dt as a function of time at 4 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 18 Table 28. Concentration of hypochlorite ion, ph, TDS, and ionic strength (μ) in OSG hypochlorite ion solutions 19 Table 29. Secondorder rate constants of perchlorate ion formation in OSG hypochlorite ion solutions, experiment vs. model (k 2obs = experimental k 2, k 2cal = predicted k 2, in units of L mol 1 day 1 ) 111 Table 3. Secondorder rate constants of perchlorate ion formation in calcium hypochlorite solutions, experiment vs. model (k 2obs = experiment k 2, k 2cal = k 2 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 viii

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

13 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.2 μg/l of perchlorate ion 13 Figure 3. Extracted ion chromatogram of perchlorate ion MRM m/z 98.9/82.8 of (a) sodium hypochlorite solution with.1 μg/l ClO 4 ; (b) standard solution with.1 μg/l ClO 4 14 Figure 4. Bromate ion chromatograms of sodium hypochlorite sample diluted by: (a) factor of 1:1; (b) factor of 1:1; (c) factor of 1:1 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 LCMS/MS, during a chlorate ion spike experiment, at 75 ºC (Control = solution at initial [OCl ] = 1.46 M and [ClO 3 ] =.29 M) 23 Figure 7. Stock solutions of ascorbic acid freshly prepared (left), after 2 days (center), and after 37 days of storage (right) 27 Figure 8. Ascorbic acid quenched hypochlorite ion sample solutions (left 3 bottles) and malonic acidquenched solution (right bottle) 28 Figure 9. Chromatogram of bromate (left) and 18 Olabeled bromate (right) of (a) sulfitequenched sample, and (b) thiosulfatequenched sample of 13% sodium hypochlorite solution diluted by a factor of 1:1, 29 Figure 1. Decomposition of hypochlorite ion and formation of chlorate ion at 75 ºC in solutions, at various initial concentrations of hypochlorite ion 4 Figure 11. Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions, at various initial concentrations of hypochlorite ion 4 Figure 12. Decomposition of hypochlorite ion and formation of chlorate ion at 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 x

14 Figure 14. Effects of Transition Metals Ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) on the hypochlorite ion decomposition and the chlorate ion formation 43 Figure 15. Effects of Transition Metals Ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) on perchlorate ion formation 44 Figure 16. Effects of noble metal ions (Noble Me = Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ), plots of hypochlorite ion decomposition and chlorite ion formation 45 Figure 17. Effects of noble metals ions (Noble Me = Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ), 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) 5 ºC, (b) 3 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) 5 ºC, (b) 3 ºC 48 Figure 2. Plots of perchlorate ion formation in solutions at various initial concentrations of chlorite ion and/or chlorate ion at (a) 3 ºC, (b) 5 ºC 49 Figure 21. Decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions spiked with bromide, chlorite, and transition metal ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) 51 Figure 22. Formation of perchlorate ion at 5 ºC in solutions spiked with bromide, chlorite, and transition metal ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) 51 Figure 23. Decomposition of hypochlorite ion at 5 ºC in solutions spiked with bromide and bromate ions at ph~ Figure 24. Formation of bromate ion at 5 ºC in solutions spiked with bromide and bromate ions at ph~ Figure 25. Formation of bromate ion at 5 ºC in solutions spiked with bromide and bromate ions at ph~ xi

15 Figure 26. Decomposition of hypochlorite ion at 4 ºC in solutions at various initial concentrations of chloride ion 56 Figure 27. Formation of chlorate ion at 4 ºC in solutions at various initial concentrations of chloride ion 57 Figure 28. Formation of perchlorate ion at 4 ºC in solutions at various initial concentrations of chloride ion 57 Figure 29. Decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in (a) 1.4 M OCl, (b).9 M OCl solutions at various initial ph 59 Figure 3. Formation of perchlorate ion at 4 ºC in (a) 1.4 M OCl, (b).9 M OCl solutions at various initial ph 6 Figure 31. Overlaid plots of (a) hypochlorite ion decomposition; (b) chlorate ion formation at 6 ºC.12 M OCl solutions at various initial ph 61 Figure 32. Overlaid plots of perchlorate ion formation at 6 ºC.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 3 ºC in solutions at various initial concentrations of chlorate ion at ph ~ Figure 34. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in solutions at various initial concentrations of chlorate ion at ph ~ Figure 35. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at various initial concentrations of chlorate ion at ph ~ Figure 36. Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 3 ºC in solutions constant in hypochlorite ion at ph ~ Figure 37. Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 4 ºC in solutions constant in hypochlorite ion at ph ~ xii

16 Figure 38. Figure 39. Figure 4. Figure 41. Figure 42. Figure 43. Figure 44. Figure 45. Figure 46. Figure 47. Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 5 ºC in solutions constant in hypochlorite ion at ph ~ Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 3 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~ Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~ Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~ Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 3 ºC in solutions at constant chlorate ion at ph ~ Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 4 ºC in solutions at constant chlorate ion at ph ~ Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 5 ºC in solutions at constant chlorate ion at ph ~ Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 75 ºC in solutions at constant chlorate ion at ph ~ Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 3 ºC in solutions at constant molar product at ph ~ Overlaid plots of (a) perchlorate ion formation; (b) decomposition xiii

17 of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at constant molar product at ph ~ Figure 48. Parallel reaction pathway linear fitted plots of (a) 3 ºC experiment; (b) 4 ºC experiment; (c) 5 ºC experiment; in solutions at constant chlorate ion and various hypochlorite ion at ph ~ Figure 49. Consecutive reaction pathway linear fitted plots of (a) 3 ºC experiment; (b) 4 ºC experiment; (c) 5 ºC experiment; in solutions at constant chlorate ion and various hypochlorite ion at ph ~ Figure 5. Rate of perchlorate ion formation as a function of (a) ionic strength; (b) concentration of hypochlorite ion at 4 ºC 82 Figure 51. Plot of μ /( 1+ μ) and bμ term as a function of ionic strength. Note: value of.5 was assumed for the b term as an approximation 86 Figure 52. Overlaid linear plots of log of secondorder rate constant versus ionic strength in solutions at various initial concentrations of hypochlorite ion at different temperatures 87 Figure 53. Smoothline 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(k /T) as a function of (1/T) 9 Figure 55. Smoothedline plots of hypochlorite ion decomposition and chlorate ion formation determined experimentally in conjunction with Bleach 21 (Error bars set at ± 1 %), for solutions at (a) [OCl ] = 82 g/l, [ClO 3 ] = 63 g/l at 3 ºC; (b) [OCl ] = 7 g/l, [ClO 3 ] = 51 g/l at 4 ºC; (c) [OCl ] = 83 g/l, [ClO 3 ] = 5 g/l at 5 ºC 94 Figure 56. Overlaid smoothedline plots of predicted (Error bars set at ± 1 %) perchlorate ion formation and determined experimentally, for solutions incubated at (a) 3 ºC; (b) 4 ºC; (c) 5 ºC 97 xiv

18 Figure 57. Overlaid smoothedline plots of (a) hypochlorite ion decomposition; (b) chlorate ion formation; (c) perchlorate ion formation determined experimentally during incubation at 5 ºC 99 Figure 58. Overlaid smoothedline plots of hypochlorite ion decomposition and chlorate ion formation determined experimentally in conjunction with Bleach 21 (Error bars set at ± 1%), for solutions with (a) [OCl ] = 63 g/l, [ClO 3 ] = 23 g/l; (b) [OCl ] = 111 g/l, [ClO 3 ] = 8.7 g/l; (c) [OCl ] = 89 g/l, [ClO 3 ] = 4.4 g/l ; (d) [OCl ] = 97 g/l, [ClO 3 ] = 12 g/l; incubated at 5 ºC 1 Figure 59. Overlaid smoothedline plots of predicted (Error bars set at ± 1 %) perchlorate ion formation and determined experimentally at 5 ºC in solutions with (a) [OCl ] = 63 g/l, [ClO 3 ] = 23 g/l; (b) [OCl ] = 111 g/l, [ClO 3 ] = 8.7 g/l; (c) [OCl ] = 89 g/l, [ClO 3 ] = 4.4 g/l; (d) [OCl ] = 97 g/l, [ClO 3 ] = 12 g/l 11 Figure 6. Smoothedline 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,.34 M ClO 3, and μ = 7.5 M; rate at 35 ºC = 1% 13 Figure 61. Overlaid smoothedline plots of (a) predicted decomposition of hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO 4 per mg OCl as a function of time in solutions at 2.3 M OCl and 1.2 M OCl at 35 ºC 15 Figure 62. Overlaid smoothedline plots of (a) predicted decomposition of hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO 4 per mg OCl as a function of time in solutions at 2.3 M OCl and 1.2 M OCl at 25 ºC 16 Figure 63. Overlaid smoothedline plots of hypochlorite ion decomposition (a) OSG solutions 16; (b) OSG solutions 712, 5 ºC 19 Figure 64. Overlaid smoothedline plots of perchlorate ion formation (a) OSG solutions 16; (b) OSG solutions 712, 5 ºC 11 Figure 65. Smoothedline plots of hypochlorite ion decomposition and xv

19 chlorate ion formation in calcium hypochlorite solutions (a) incubated at 5 ºC; (b) incubated at 6 ºC 113 Figure 66. Smoothedline plots of perchlorate ion formation in calcium hypochlorite solutions (a) incubated at 5 ºC; (b) incubated at 6 ºC 113 Figure A11. SEM images of typical gold nanoislands produced by sputtering process on a polished aluminum substrate 124 Figure A12. AFM images of 25 nm (a); and 14 nm (b) gold nanoislands on quartz substrate 125 Figure A13. Overlaid UVVis spectra of : dashed linegold thin film on quarts, solid linegold nanoislands, dotdashed line gold nanoislands exposed to ozone gas, dotted linegold nanoislands reversed by annealing at 375 ºC for 15 min 126 Figure A14. Overlaid UVVis spectra of 25 nm gold nanoislands with surface Plasmon absorbance max at 52 nm exposed to concentrations of ozone, increased in increments form 2.9 μg/l to 166.1μg/L. Ozone causes a redshift in the surfaceplasmon absorbance max 128 Figure A15. Shift of the 25 nm gold nanoislands surfaceplasmon max (52 nm) as a function of ozone concentration, logarithmic fit gives an equation of y=6.8ln(x)15.19 produced a correlation coefficient of Figure A16. Surface plasmon s shifts of gold nanoislands with absorbance max At 532 nm as a function of ozone concentration 13 Figure A21. (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 A22. EPONCoated ITO Electrode, plating time 1 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 xvi

20 Figure A23. EPONcoated ITO electrode, plating time: (a) 3 s 2.1V. The Deposited film indicates normal to the electrode surface orientation of the deposited EPON film. (b) EPONcoated ITO electrode, magnification of (a) reveals normal orientation of the mesopores 139 Figure A24. (a) Dry SolGel 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 4methylcoumarin7yl 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 4methylcoumarin7yl 3(trimethoxysilyl)propylcarbamate, that is not bound; dried and wetted again 14 Figure A25. Overlaid excitation spectra of EPON film subjected to different ph: (a) DI water; (b) ph 1; (c) ph 2.2; (d) ph Excitation maximum and excitation peak shape shifts based on ph of the wetting solution 141 Figure A26. 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 xvii

21 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 xviii

22 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 ontrack. 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 solgel and electrochemical methods (Appendix 1). I would like to thank Dr. Diep Ca for xix

23 help with ICPOES 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 ICPMS, and for the training to perform analysis by MALDITOF and LCMS/MS. I wish to acknowledge Barry Landrum for the undisputedlysuperb machinist skills to craft specialtydesigned 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 TaulbeeCombs, 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. xx

24 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 29, American Water Works Association, cosponsored by AWWA Water Industry Technical Action Fund (WITAF), the Water Research Foundation, and the Southern Nevada Water Authority. xxi

25 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 reductionoxidation behavior of the perchlorate ion is rarely observed and because perchlorate ion has less Figure 1. Perchlorate ion 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 ClO 4 + 8H + + 8e = Cl + 4H 2 O E o = V at ph (1) 1 CDHS, "Determination of Perchlorate by Ion Chromatography", California Department of Health Services (CDHS), 1997, Retrieved from: 2 CDPH, "History of Perchlorate in California Drinking Water", California Department of Public Health (CDPH), 27, Retrieved 7/28/9, from: 3 Robinson, R. A. and Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths Publications Limited: London, Housecroft, C. E. and Sharpe, A. G. Inorganic Chemistry; Pearson Education Limited, 21. 1

26 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) light 9 and in the presence of atmospheric ozone. 1 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, tons in the 195s to the peak of 15, ton in mid198s. 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 25, 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. 21, 112, 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 (Cl 4 ), EPA, Urbansky, E. T. "Perchlorate as an environmental contaminant" Env. Sci. Pollut. Res. 22, 9, Kang, N., Anderson, T. A. and Andrew Jackson, W. "Photochemical formation of perchlorate from aqueous oxychlorine anions" Anal. Chim. Act. 26, 567, 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. 28, 45, Mack, E. L. "Electrolytic Formation of Perchlorate" J. Phys. Chem. 1917, 21, Davis, C. O. Perchlorate explosives, E. I. du Pont de Nemours & Co., 194, USA US Motzer, W. "Perchlorate: Problems, Detection, and Solutions" Environ. Forensics 21, 2,

27 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 25 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 longterm environmental concern. 1.2 Perchlorate Ion: Toxicity and Regulation Perchlorate ion is an endocrinedisrupting 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 25 the National Academy of Sciences and the United States Environmental Protection Agency (USEPA) established an RfD of.7 mg/kg/day, which corresponds to a drinking water equivalent level (DWEL) of 24.5 µg/l. 2 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, 28 an interim health advisory level of 15 μg/l (based on 14 USEPA, "Known Perchlorate Releases in the U.S. March 25th, 25", United States Environmental Protection Agency, 25, Retrieved 8/1/9, from: 15 Tikkanen, M. W. "Development of a drinking water regulation for perchlorate in California" Anal. Chim. Act. 26, 567, USGAO Perchlorate: A System to Track Sampling and Cleanup Results Is Needed United States Government Accountability Office (USGAO): Washington, D.C., 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. 26, 567, 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. 22, 11, 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. 24, 39, USEPA Draft Regulatory Determinations Support Document for Selected Contaminants from the Second Drinking Water Contaminant Candidate List (CCL2) EPA Report 815D67, United States Environmental Protection Agency (USEPA): Washington, DC., USFDA, "2425 Exploratory Survey Data on Perchlorate in Food", United States Food & Drug Administration, 25, Retrieved 8/2/9, from: ucm77685.htm 3

28 the RfD of.7 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 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 27 survey, 63 % of water treatment facilities use chlorine as the primary disinfectant. As many as 3% of the drinking water treatment facilities (DWTF) in North America had switched from using chlorine gas to hypochlorite ion solutions in the past 1 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 onsite 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, 28, Retrieved 7/27/9, from: 23 Greiner, P., Mclellan, C., Bennet, D. and Ewing, A. "Occurrence of perchlorate in sodium hypochlorite" J. Am. Water Work Assoc. 28, 1, 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. 29, 58, Renner, R. "EPA perchlorate decision flawed, say advisers" Environ. Sci. Technol. 29, 43, Shah, J. and Qureshi, N. "Chlorine Gas vs. Sodium Hypochlorite: What's the Best Option" Opflow 28, Routt, J., Mackey, E. and Noack, R. "Committee Report: Disinfection Survey, Part 2 Alternatives, experiences, and future plans" J. Am. Water Work Assoc. 28, 1, 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," 24, Retrieved 4/14/9, from: PotableWaterTreatment.pdf 4

29 hypochlorite ion itself, and chlorite ion. 3 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 solutions 36 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 chloralkali 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 3 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, Weinberg, H. S., Delcomyn, C. A. and Unnam, V. "Bromate in Chlorinated Drinking Waters: Occurrence and Implications for Future Regulation" Environ. Sci. Technol. 23, 37, Bouland, S., Duguet, J. P. and Montiel, A. "Evaluation of bromate ions level introduced by sodium hypochlorite during postdisinfection of drinking water" Environ. Technol. 25, 26, 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," 26, Retrieved 7/8/9, from: 38 Adam, L. C. An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite Ph.D. Dissertation, Miami University, Oxford, OH, Adam, L. C. and Gordon, G. "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion" Inorg. Chem. 1999, 38,

30 decomposition model was developed (referred as Bleach 21 in later chapters). 4 As a result of these studies, it was found that hypochlorite ion solutions are most stable in the ph range of 1113 and at lower temperatures. Decomposition of hypochlorite ion, in alkaline solutions, is secondorder 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 ± 1 %. vi. Determine strategies to minimize perchlorate ion formation in concentrated sodium hypochlorite solutions. Develop recommendations to water utilities that use hypochlorite ion for disinfection. 4 Adam, L., Gordon, G. and Pierce, D., Bleach 21 Predictive Model, Miami University, Oxford, Ohio. Copyright (c) 21 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. 6

31 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 LCMS/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 LCMS/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 LCMS/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 ionchromatography (ICMS) 43, 44 or liquid chromatography (LCMS) 45,46 has been utilized for the detection and quantification of oxyhalide anions. Both US EPA method 331., that uses LC with tandem mass spectrometry (LCMS/MS), 47 and US EPA method 332., 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 2, 888, 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 22, 956, 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, Urbansky, E. T., Magnuson, M. L., Freeman, D. and Jelks, C. "Quantitation of perchlorate ion by electrospray ionization mass spectrometry (ESIMS) using stable association complexes wih organic cations and bases to enhance selectivity" J. Anal. At. Spectrom. 1999, 14, Wendelken S. C., V. L. E., Coleman D. E., Munch D. J. U.S. Environmental Protection Agency Method 331. Revision

32 MS/MS 48 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.35 13%, as active chlorine, with specific conductance ranging from 51.8 ms/cm to 498 ms/cm (ionic strength ranging from.88m). 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 LCMS/MS in water 49, which may necessitate higher dilutions or sample cleanup 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 solution 5, quantitation of such analytes must also be validated. Recently, it has been demonstrated that LCMS/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 water 51 and food supplements 52. However, these methods are generally timeconsuming 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 titrations 54 have 48 Hedrick, E. B., T.;Slingsby, R.;Munch, D. U.S. Environmental Agency Method 332. Revison Li, Y. E., J. George "Analysis of Perchlorate in Water by ReversedPhase LC/ESIMS/MS Using an Internal Standard Technique" Anal. Chem. 25, 77, 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. 25, 39, 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. 26, 567, Clesceri, L. S. G., A. E.; Eaton, A. D. Eds. Standard Methods for the Examination of Water and Wastewater 2th ed.; American Public Health Association: Washington, DC,

33 been used, but are too timeconsuming 55 or are optimized for determination of a single oxyhalide species 56. 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 Transition metal ions: Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ Metal ion analysis and quantification was performed using US EPA method on an Agilent (Palo Alto, CA) 75c Inductively Coupled PlasmaMass Spectrometer (ICPMS) with an Octopole Reaction System that uses hydrogenhelium reaction gas to remove Arbased isobaric and polyatomic oxide interferences. Internal standards were used to correct for matrix interferences. Dilutions ranging from 1:1 to 1:5 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. ICPMS MRLs (µg/l) in water and hypochlorite ion solutions Element Reagent Water Hypochlorite Ion Solutions* Manganese (Mn) Iron (Fe) Cobalt (Co) Copper (Cu) Nickel (Ni) *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, Ikeda, Y., Tang, T. and Gordon, G. "Iodometric method for determination of trace chlorate ion" Anal. Chem. 1984, 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, Creed, J. T. B., C.A.; Martin, T.D. U.S. Environmental Agency Method 2.8 Revison

34 2.1.3 Specific Conductance, Ionic Strength, and ph Measurements Specific conductance measurements were performed using a HACH IonSeries Conductivity / Total Dissolved Solids meter (Hach Company, Loveland, CO). A calibration was performed using a 1, 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 22, 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 1 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 1 4 (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, 1). 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 The LCMS/MS Analysis of Perchlorate, Bromate, and Chlorate Ions Tandem mass spectrometry was performed using an API4 triplequadrupole 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 collisioninduced 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 1

35 MS/MS very selective and sensitive 58. The following precursor/product ion transitions were used: 35 ClO 4 (m/z 99) to 35 ClO 3 (m/z 83) for perchlorate ion; 35 ClO 3 (m/z 83) to 35 ClO 2 (m/z 67) for chlorate ion; and 79 BrO 3 (m/z 127) to 79 BrO 2 (m/z 111) for bromate ion. The following confirmation transitions were used: 37 ClO 4 (m/z 11) to 37 ClO 3 (m/z 85) for perchlorate ion; 37 ClO 3 (m/z 85) to 37 ClO 2 (m/z 69) for chlorate ion; and 81 BrO 3 (m/z 129) to 81 BrO 2 (m/z 113) for bromate ion. Perchlorate and bromate ions were quantified using isotope dilution. Stable isotopelabeled versions of perchlorate ( 35 Cl 18 O 4 ) and bromate ( 79 Br 18 O 4 ) ions were used. As no source of oxygen18 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 stableisotope labeled perchlorate (Cl 18 O 4 ) and bromate (Br 18 O 3 ) 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 MilliQ Gradient System (Millipore, Billerica, MA). In order to reduce the run times, analytes were separated using a 75 x 4.6 mm Synergi MaxRP C12 column with a 4 μm pore size (Phenomenex, Torrance, CA). An injection volume of 2 μl was used for all samples. A binary gradient consisting of.1 % formic acid (v/v) in water (A) and 1 % methanol (B) at a flow rate of 7 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,

36 original method 59 with all three analytes eluting after two minutes. The effects of faster separation on method performance were further investigated Validation of LCMS/MS Method for the Analysis of Hypochlorite Solutions Using the optimized separation method, standards containing.2,.1, and.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.2 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 1%) of 8 replicate measurements of the.2 μg/l perchlorate ion standard is almost twice as large as that of the.1 μg/l perchlorate ion standard but is still below 1% RSD. Bromate and chlorate ions MDLs of.6 μg/l and.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.15 μg/l was used for ClO 4 ). This resulted in PQLs of.5 μg/l for ClO 4,.2 μg/l for BrO 3, and.7 μg/l for ClO 3. Analysis of most sodium hypochlorite solutions for measurement of perchlorate ion rarely required a PQL<.1 μg/l, thus the lowest calibration standard for typical analysis contained.1 μg/l of ClO 4,.2 μg/l of BrO 3, and 1. μg/l of ClO 3. To demonstrate the applicability of the PQL to a complex sample matrix, chromatograms of sample and standard perchlorate ion solutions are shown in Figure Snyder, Vanderford and Rexing "Trace analysis of bromate, chlorate, iodate, and perchlorate in natural and bottled waters". 12

37 Table 2. MDL data for perchlorate, bromate, and chlorate ions (n = 8) Standard Concentration Replicate.2 μg/l.1 μg /L.1 μg /L.5 μg /L Mean Std. Dev RSD (%) Student' 98% n MDL = SD X Student's t PQL= 3 x MDL(μg /L)* *Values rounded up Intensity, cps μg/l ClO 4 Standard Solution 2.3 S/N = 6.3 Peak Int.(Subt.)=7.e+2 3xStd.Dev.(Noise)=1.1e Time, min Figure 2. Extracted ion chromatogram of perchlorate ion MRM m/z 98.9/82.8 of standard solution containing.2 μg/l of perchlorate ion 13

38 (a) Intensity, cps μg/l ClO 4 Sodium Hypochlorite Solution 2.1 S/N = 13.5 Peak Int.(Subt.)=2.2e+3 3xStd.Dev.(Noise)=1.6e Time, min (b) Intensity, cps μg/l ClO 4 Standard Solution 2.3 S/N = 12.9 Peak Int.(Subt.)=2.6e+3 3xStd.Dev.(Noise)=2.e Time, min Figure 3. Extracted ion chromatogram of perchlorate ion MRM m/z 98.9/82.8 of (a) sodium hypochlorite solution with.1 μg/l ClO 4 ; (b) standard solution with.1 μg/l ClO 4 The signaltonoise (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.1 μg/l ClO 4, 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/x 2 14

39 weighting and typical R For chlorate ion, an external calibration with 1/x 2 weighting was used with typical R 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 LCMS/MS. Due to known isobaric interferences by bisulfate ion (HSO 4 ) 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 3 ml deionized water. Sample solutions were eluted typically at a flow rate of 2. ml/min using a mechanical syringe pump (KD Scientific, Holliston, MA) and at least the first 1 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 cleanup 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 cleanup/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) ClO 4 BrO 3 ClO 3 No Filtration Ba/H Filtration Dilution Factor (DF) % Recovery * 91 Std. Dev RSD (%) % Recovery * 98 Std. Dev RSD (%) MRL =PQL X DF (μg/l) , ,5 *Dilution is inadequate resulting in concentrations impeding analysis 15

40 The analysis of perchlorate ion showed recoveries of 99.5 % (± 3.6) for a wide range of dilutions, analyte concentrations, and Ba/H cleanup steps. Perchlorate ion could be quantified at dilutions as low as 1:1, thereby achieving MRL as low as.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. (a) 79 Br 16 O 3 79 Br 18 O 3 (b) 79 Br 16 O 3 79 Br 18 O 3 (c) 79 Br 16 O 3 79 Br 18 O 3 Figure 4. Bromate ion chromatograms of sodium hypochlorite sample diluted by: (a) factor of 1:1; (b) factor of 1:1; (c) factor of 1:1 16

41 Interestingly, the loss of 18 Olabeled bromate ion signal was much higher than that of the analyte signal at 1:1 and 1:1 dilutions, thereby producing erroneous results (figure 3). A higher dilution of 1:1 eliminated this problem since the peakwidening 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 LCMS/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. mg/l ClO 4 and BrO Bromate Chlorate Perchlorate mg/l ClO 3 Figure Dilution Factor 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. 17

42 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 (I 3 ), in presence of excess iodide ion, shown by Equations 46: OCl + 3I + 2H + I 3 + Cl + H 2 O ph 1.3 (4) ClO 2 + 6I + 4H + 2I 3 + Cl + 2H 2 O ph 1.3 (5) ClO 3 + 9I + 6H + 3I 3 + Cl + 6H 2 O 6M H + (6) Triiodide ion can be titrated with thiosulfate ion (S 2 O 2 3 ) or sulfite ion (SO 2 3 ), shown by Equations 7 and 8, thus allowing determination of these oxyhalide anions. I S 2 O 3 3I 2 + S 4 O 6 (7) I 3 + SO OH SO I + H 2 O (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 (IO 3 ) to an acidic solution containing a small excess of iodide ion, shown in Equation 1. Thus, the prepared standard solution of triidodie ion can be used to standardize thiosulfate ion and sulfite ion solutions. IO 3 + 8I + 6H + 3I 3 + 3H 2 O (9) Similarly, bromate ion reacts with iodide ion, shown by Equation 1. BrO I + 6 H + 3 I H 2 O ph 1.3 (1) 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. 18

43 AdamGordon Method Direct potentiometric titration of hypochlorite ion with sulfite ion allows selective determination of hypochlorite ion from chlorite ion and bromate ions. 6 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 1.5, using.4 M borate buffer (b) Titrate hypochlorite ion with sulfite ion, given by equation 11: OCl 2 + SO 3 SO Cl ph 1.5 (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: I 3 + SO OH SO I + H 2 O (9) (d) Remove any excess triiodide ion by adding dilute thiosulfate ion solution: I S 2 O 3 3I 2 + S 4 O 6 (8) Step 2: Determination of chlorite ion: a) Adjust ph to 1.3 using 3 M HCl, reaction 4 takes place: ClO 2 + 6I + 4H + 2I 3 + Cl + 2H 2 O (5) b) Backtitrate triiodie ion with thiosulfite ion: I S 2 O 3 3I 2 + S 4 O 6 (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: ClO 3 + 9I + 6H + 3I 3 + Cl + 3H 2 O 6M H + (6) b) Backtitrate triiodie ion with thiosulfite ion: I S 2 O 3 3I 2 + S 4 O 6 (8) 6 Adam and Gordon "Direct and Sequential Potentiometric Determination of Hypochlorite, Chlorite, and Chlorate Ions When Hypochlorite Ion is Present in Large Excess". 19

44 Titrations with sulfite ion were performed using a VIT 9 Video Titrator with a P11 platinum k41 SCE electrode pair (Radiometer, Copenhagen, Denmark). Titrations with thiosulfate ion were performed with a standard 5 ml laboratory glass burette. Borate buffer (.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 1.5. 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 ( % 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 SigmaAldrich (St. Louis, MO). Dilute ( 1 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 (.1 M) was prepared weekly and used to standardize sulfite and thiosulfate ion solutions daily. A solution containing ~.2 M I 3 was prepared by adding an aliquot of ~.1 M IO 3 standard to acidic solution (ph ~1) containing excess iodide ion (~.5 M I ) and used to remove 2 sulfite ion. Standardization of.2 M SO 3 and.1 M S 2 O 2 3 solutions by.1 M IO 3 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.19 M IO 3 ml of IO 3 Standard 2 ml SO 3 [SO 2 3 ] ml of IO 3 Standard 2 ml S 2 O 3 [S 2 O 2 3 ] Mean.241 Mean.77 5 Std. Dev.. 9 Std. Dev.. 5 RSD (%).353 RSD (%).685 2

45 2.2.4 Method Selection for the Measurement of Chlorate Ion Sodium hypochlorite solutions are obtained as either bulk commercial bleach as 313 % free available chlorine (FAC), or as onsite generated (OSG) sodium hypochlorite solutions that can range.3 3 % as FAC. Bulk sodium hypochlorite solutions, 1 13% FAC, were obtained from Acros Organics USA (Morris Plains, NJ) and VWR (Brisbane, CA). OSG sodium hypochlorite solution, ~.8 % FAC, was obtained from a sodium hypochlorite onsite generator at River Mountains Water Treatment Facility (RMTF) (Henderson, NV). To determine which method (iodometric titration or LCMS/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 LCMS/MS and iodometric titration for bulk hypochlorite ion solutions (n = 7) Nonspiked Hypochlorite ClO 3 Spiked % Recovery Replicate LCMS/MS* Titration* LCMS/MS* Titration* LCMS/MS* Titration* Mean Std. Dev RSD (%) *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 LCMS/MS. This higher variability observed in the LCMS/MS data is most likely a result of carrying out serial dilutions of several orders of 21

46 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 1 g/l, a dilution factor of at least 4, would be needed to bring the concentration of chlorate ion to the middle of a typical calibration curve of 1 5 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 1:2,5, were required. The error bars of chlorate ion measurements by LCMS/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 ±.2 g/l (Table 5) or ±.2 M, the iodometric titration method should be used to measure concentration of chlorate ion above 2. g/l (SD=.2 g/l X 1) or.25 M to achieve reasonable precision (RSD < 1 %). The results shown in Table 6 indicate poor precision when measuring concentration of chlorate ion less than 1. g/l by iodometric titration. Table 6. Comparison of measurements by the LCMS/MS and iodometric titration for OSG sodium hypochlorite solutions at less than 1. g/l ClO 3 (< 1 mm) (n 3) Titration LCMS/MS Mean (g/l ClO 3 ) Std. Dev. (g/l ClO 3 ).2.1 RSD (%) (n=4) (n=3) 22

47 (a) (b) g/l ClO g/l ClO Control Control + 5g/L ClO 3 Control + 1g/L ClO 3 Control + 15g/L ClO Days Control Control + 5g/L ClO 3 Control + 1g/L ClO 3 Control + 15g/L ClO 3 Figure Days Comparison of chlorate ion measurements by (a) iodometric titration and (b) by the LCMS/MS, during a chlorate ion spike experiment, at 75 ºC (Control = solution at initial [OCl ] = 1.46 M and [ClO 3 ] =.29 M) Although accurate and reproducible trace analysis of chlorate ion by iodometric titration is possible 61, 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 1 25 g/l range (generally 61 Ikeda, Tang and Gordon "Iodometric method for determination of trace chlorate ion". 23

48 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 1:1,, chlorate ion measurements by LCMS/MS method, typically had an RSD of less than 5 % (n = 3). Thus the LCMS/MS method can be used reliably for measurement of chlorate ion in the.1 1, mg/l range, but was not required for chlorate ion spike experiments 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 13 days; at 4 ºC the half life is 3184 days, according to the Bleach 21. 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 Wood 62, Sweetin 63, and Adam 64 at Miami University. For example, ascorbic acid was found to be most effective over the ph range of 311, glycine above ph 8.5, oxalic acid below ph 8.5, malonic acid over the ph range of A selective reaction of sulfite ion and hypochlorite ion at ph 1.5 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, 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, Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 24

49 development of an analytical method for the accurate determination of hypochlorite ion 65, Equation 11: OCl 2 + SO 3 SO Cl ph 1.5 (11) At ph 1.5 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: S 2 O HOCl + H 2 O 2SO Cl + 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 + H 2 O 2 Cl + H 2 O + O 2 (13) Hydrogen peroxide is also known to react with iodide ion 66 as well as sodium thiosulfate and sodium sulfite, 67 as given by Equations Thus, any excess hydrogen peroxide from the reaction with hypochlorite ion can potentially interfere with the measurement of chlorite and chlorate ions. 3I + H 2 O 2 + 2H + I 3 + 2H 2 O (14) 2S 2 O H 2 O 2 S 4 O OH (15) SO H 2 O 2 SO H 2 O (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 (LCMS/MS method) Impact on bromate ion levels (LCMS/MS method) Impact on chlorate ion levels (LCMS/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 Thirdorder Ionic Reaction without Appreciable Salt Effect" J. Phys. Chem. 1934, 38, Liu, W., Andrews, S. A., Stefan, M. I., & Bolton, J. R. "Optimal methods for quenching H 2 O 2 residuals prior to UFC testing" Water Res. 23, 37,

50 Ascorbic acid, USP Grade, was obtained from Mallinckrodt Chemicals (Philipsburg, NJ). Glycine, ACS Reagent 98.5%, was supplied by Sigma Aldrich. Hydrogen peroxide, 332%, Assay % 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 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 preweighed (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 1 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 H 2 O 2. Quenching 1 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 wellventilated chemical hood, the gas produced from reaction between glycine and hypochlorite ion caused feelings of lightheadedness, 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 easiertohandle quenching agents tested. None of the quenching agents had associated 26

51 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 shelflife. 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 2 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 2 days (center), and after 37 days of storage (right) 27

52 Figure 8. Ascorbic acid quenched hypochlorite ion sample solutions (left 3 bottles) and malonic acidquenched solution (right bottle) 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 predissolved 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,.75 for malonic acid, and 1.5 for oxalic acid. Glycine was the least reproducible in quenching hypochlorite ion with mole ratios varying from.2 to.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. 28

53 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 LCMS/MS. Both bromate ion and the bromate ion internal standard were negatively impacted by the presence of thiosulfate and sulfite ions. The chromatograms of sulfitequenched and thiosulfatequenched samples, shown in Figure 9, illustrate the peak shifts, peak attenuation, and peak splitting (Figure 4 (c), page 16 can be used as reference). (a) 79 Br 16 O 3 79 Br 18 O 3 (b) 79 Br 16 O 3 79 Br 18 O 3 Figure 9. Chromatogram of bromate (left) and 18 Olabeled bromate (right) of (a) sulfitequenched sample, and (b) thiosulfatequenched sample of 13% sodium hypochlorite solution diluted by a factor of 1:1, Furthermore, suppression of both analyte and internal standard signals hindered the ability for adequate correction by isotopedilution, resulting in poor recoveries for samples quenched with thiosulfate ion (6%, n = 3) and no quantifiable recoveries for 29

54 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 acidquenched hypochlorite ion solutions with bromate ion standard, showed much lower recoveries (49%, n=3) than the nonquenched 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 nonquenched hypochlorite ion samples are shown in Table 7, while comparisons of quenched and nonquenched 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 LCMS/MS (n = 3, replicate samples analyzed in triplicate; S.D. = standard deviation) % Recovery (Mean ± S.D.) Perchlorate Spike μg/l Nonquenched sample MA quenched sample ± ± ± ± 1.2 Bromate Spike μg/l ± ± 1.1 Chlorate Spike μg/l ± ±.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 ClO 4 mg/l ± S.D. Sample BrO 3 mg/l ± S.D. Dilution Nonquenched sample MA quenched sample Nonquenched sample MA quenched sample 1:1, 6.9 ± ± ± ±.9 1:1, 7.2 ± ± ± ±.1 1:5 7.4 ± ± ± ±.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 26, 221,

55 Table 9 shows measured concentrations of chlorate ion by LCMS/MS and iodometric titration in nonquenched sodium hypochlorite and in malonic acid (MA) quenched sample. The difference in results and recoveries between quenched and nonquenched 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 ClO 3 g/l ± S.D. Nonquenched Hypochlorite Malonic Acid LCMS/MS 16.1 ± ±.15 Iodometric Titration 18.1 ± ±.2 % Difference 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 1% difference on average (based on 8 different hypochlorite ion samples) in concentration of perchlorate, bromate, and chlorate ions after 2month storage, shown in Tables 112. Table 1. Perchlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period μg/l ClO 4 Sample Index 4/2/29 6/2/29 % Difference Mean 7. ± 5. 31

56 Table 11. Bromate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period mg/l BrO 3 Sample Index 4/2/29 6/2/29 % Difference Mean 1. ± 5.3 Table 12. Chlorate ion stability in hypochlorite ion solutions quenched with malonic acid over 2month period mg/l ClO 3 Sample Index 4/2/29 6/2/29 % Difference Mean 9.7 ± 5. 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 Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ by ICPMS method against the nonquenched sample. No significant differences were observed, as most of metals ions were already present at very low concentrations 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 32

57 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 offsite and requiring shipping). Table 13. Stoichiometric Ratio (mol QA / mol OCl ) Summary of quenching agent test results and decisionmaking matrix Ascorbic Acid Glycine Malonic Acid Quenching Agent (QA) Sodium Thiosulfate Sodium Sulfite Oxalic Acid Hydrogen Peroxide Compatible with Iodometric Titration? Effect on [BrO 3 ] No Yes Yes Yes Yes Yes No Decrease No Effect No Effect Decrease Decrease No Effect No Effect Effect on [ClO 4 ] No Effect No Effect No Effect No Effect No Effect No Effect No Effect Effect on [ClO 3 ] No Change No Change No Change No Change No Change No Change No Change Solution Stable over Time Safety Issues? Soluble in Concentrated NaOCl? No Yes Yes Yes No Yes No No Noxious Gas, Hazard! No No No Heat, Gas Evolved Violent Reaction, Hazard! Yes Yes Yes Yes Yes No Yes 33

58 2.3 Conclusions The optimized LCMS/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.5 μg/l for ClO 4 and.2 μg/l for BrO 3. The advantage of this method is twofold: analysis time of 1 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 18 Olabeled 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 18 Olabeled 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 LCMS/MS method was used for chlorate ion analysis for samples with 1 g/l ClO 3, 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 LCMS/MS method PQL of.7 μg/l (8 nm), and iodometric titration determinations of concentrations of chlorate ion up to 21 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 LCMS/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 cleanup/filtration step with Barium Cartridges to remove the interference with the perchlorate ion determination, does not negatively impact analysis by the LCMS/MS method. 34

59 The sodium hypochlorite samples, pending analysis, were temporary stored at 4 ºC (typically samples were analyzed within oneweek period). Alternatively, malonic acid was used during collection of hypochlorite ion solutions offsite, and water samples that contained residual FAC. Thus, a robust, sensitive LCMS/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. 35

60 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. % FAC) sodium hypochlorite solutions and to increase over time 69 and the fact that aged sodium hypochlorite solutions containing high concentrations of chlorate ion also had higher perchlorate ion concentrations 7, 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 313 % as FAC and have ph above The onsite generated (OSG) sodium hypochlorite solutions range.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 range, is secondorder in hypochlorite ion and is at a minimum in the ph 1213 range 71. Decomposition of hypochlorite ion is acidcatalyzed and becomes more rapid at lower ph. In the ph region of 58, decomposition of hypochlorite ion is a thirdorder process. 72 Given, that the pk a of hypochlorous acid 73 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". 7 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 58 Region" Inorg. Chem. 1992, 31, 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,

61 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 Cu 2+ and Ni 2+ were specifically shown to catalyze decomposition. 76 Furthermore, iridium ion (Ir 3+ ) 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 (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) and noble metal ions (Ag +, Au +, Ir 3+ Pd 2+, and Pt 2+ ). 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.5 % of hypochlorite ion concentration). 79 OCl + OCl ClO 2 + Cl k ClO2 (slow) (17) OCl + ClO 2 ClO 3 + Cl k ClO3 (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, 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, Lister, M. W. "Decomposition of Sodium Hypochlorite: The Uncatalyzed Reaction" Can. J. Chem. 1956, 34, Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 37

62 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 (BrO 3 ), a known contaminant found in sodium hypochlorite solutions, 8 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 nonspiked hypochlorite ion solutions were stored to allow decomposition of hypochlorite ion to occur. As an aid in planning the incubation experiments, Bleach 21 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 (1114), and sodium chloride concentrations. Because the decomposition of ph hypochlorite ion solutions at room temperature is relatively slow (the half life of 13 % FAC solution at 25 ºC is 13 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, 6 ºC, and 5 ºC. Bulk sodium hypochlorite solutions, 1 13% FAC, were obtained from Acros Organics USA (Morris Plains, NJ) and VWR (Brisbane, CA). OSG sodium hypochlorite solution, ~.8 % FAC, was obtained from sodium hypochlorite onsite 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, 8 Chlorine Institute, "Bromate in Sodium Hypochlorite Potable Water Treatment" 38

63 8%) 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, ppm Co 2+, Cu 2+, Fe 3+, Ni 2+, and Mn 2+ 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, ppm Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ 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, generalpurpose, 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 acidwashed, 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 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 1 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, 39

64 chlorate, and perchlorate ions for duplicate samples were monitored during ten days of incubation at 75 ºC in a water bath. Figure 1 shows overlaid smoothedline plots of hypochlorite ion decomposition and chlorate ion formation. Note: the downward sloping smoothed lines represent hypochlorite ion decomposition, and the upwardsloping curves show formation of chlorate ion shown in Figure 1. As expected, solutions with the higher hypochlorite ion concentrations decompose more rapidly, and form higher concentrations of chlorate ion, as depicted in Figure 1. Figure 11 shows a smoothedline plot of perchlorate ion formation for this set of samples OCl g/l Decomposed g/L OCl + 5 g/l ClO 3 5g/L OCl + 5 g/l ClO 3 1g/L OCl + 5 g/l ClO ClO 3 g/l Formed Days Figure 1. Decomposition of hypochlorite ion and formation of chlorate ion at 75 ºC in solutions, at various initial concentrations of hypochlorite ion ClO 4 mg/l Formed g/L OCl + 5 g/l ClO 3 5g/L OCl + 5 g/l ClO 3 1g/L OCl + 5 g/l ClO 3 2 Figure Days Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions, at various initial concentrations of hypochlorite ion 4

65 A sample, with an initial concentration of 74 g/l OCl produced significantly higher concentration of perchlorate ion than samples with 5 g/l and 1 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 1 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 15 g/l was added. Figure 12 shows smoothedline plots of hypochlorite ion decomposition and formation of chlorate ion. Similar to Figure 1, decomposition of hypochlorite ion at 75 ºC is rapid. Figure 13 shows a smoothedline 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. 41

66 OCl g/l Decomposed g/l OCl + 24 g/l ClO 3 74 g/l OCl + 74 g/l ClO 3 72 g/l OCl g/l ClO 3 7 g/l OCl g/l ClO Days ClO 3 g/l Formed Figure 12. Decomposition of hypochlorite ion and formation of chlorate ion at 75 ºC in solutions, at various initial concentrations of chlorate ion ClO 4 mg/l Formed g/l OCl + 24 g/l ClO 3 74 g/l OCl + 74 g/l ClO 3 72 g/l OCl g/l ClO 3 7 g/l OCl g/l ClO 3 25 Figure Days Formation of perchlorate ion at 75 ºC in hypochlorite ion solutions, at various initial concentrations of chlorate ion 42

67 3.4 Effect of Transition Metal Ions (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) To investigate the effects of transition metal ions, aliquots of 1, ppm Co 2+, Cu 2+, Fe 3+, Ni 2+, and Mn 2+ standards (SPEX CertiPrep, Inc., Metuchen, NJ) were spiked into sodium hypochlorite solutions to achieve final concentration of 2 mg/l, 2 mg/l, and.2 mg/l of each metal ion. Following addition of a 2 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 3 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 Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+, 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. OCl g/l Decomposed g/l OCl + 24 g/l ClO 3 (Control) 75 g/l OCl + 24 g/l ClO mg/l Me 75 g/l OCl + 24 g/l ClO mg/l Me Days ClO 3 g/l Formed Figure 14. Effects of Transition Metals Ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) on the hypochlorite ion decomposition and the chlorate ion formation Observations, drawn from Figures 14 and 15, suggest that the presence of Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ even at a concentration of.2 mg/l exhibits a much stronger effect on the decomposition of hypochlorite ion such that no additional perchlorate ion 43

68 formation was observed. Thus, these metal ions can not enhance the rate of perchlorate ion formation under these conditions (ph 1113, FAC > 5. %), because the primary reactant to produce perchlorate ion (i.e. hypochlorite ion) is decomposed rapidly. ClO 4 mg/l Formed g/l OCl + 24 g/l ClO 3 (Control) 75 g/l OCl + 24 g/l ClO mg/l Me 75 g/l OCl + 24 g/l ClO mg/l Me Days Figure 15. Effects of Transition Metals Ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) on perchlorate ion formation 3.5 Effect of Noble Metal Ions (Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ) Similar to the investigation of transition metal ion effects on the rate of perchlorate ion formation, aliquots of 1, ppm Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ standards (Elements Inc., Shasta Lake, CA) were spiked into sodium hypochlorite solutions. The final concentration of each metal ion was.2 mg/l. Spiked samples were divided into two groups: a set with a spike of.2 mg/l Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+, and a set with a spike of.2 mg/l Ir 3+ only. This was done to separate the effects of Ir 3+, 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 5 ºC, and the target analytes were monitored. Figure 16 shows overlaid smoothedline plots of hypochlorite ion decomposition and formation of chlorate ion. 81 Ayers "Catalytic Decomposition of Hypochlorite Solution by Iridium Compounds. II. Kinetic Studies". 44

69 OCl g/l Decomposed No Spike (Control: 73 g/l OCl + 27 g/l ClO 3 ) Set 1:.2 mg/l Noble Me Ions 6 Set 2:.2 mg/l Iridium Ion Days Figure 16. Effects of noble metal ions (Noble Me = Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ), plots of hypochlorite ion decomposition and chlorite ion formation ClO 3 g/l Formed It is evident from Figure 16 that at.2 mg/l spike, unlike the transition metal ions, the noble metal ions (Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ) 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. ClO 4 mg/l Formed No Spike (Control: 73 g/l OCl + 27 g/l ClO 3 ) Set 1:.2 mg/l Noble Me Ions Set 2:.2 mg/l Iridium Ion Days Figure 17. Effects of noble metals ions (Noble Me = Ag +, Au +, Ir 3+, Pd 2+, and Pt 2+ ), overlaid plots of perchlorate ion formation 45

70 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.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 +, Ir 3+, Pd 2+, and Pt 2+ (Noble Me) vs. control (no spike), incubated at 5 ºC ClO 4 (mg/l) Day Sample Index: Set 1:.2 mg/l Noble Me Set 1:.2 mg/l Noble Me (Duplicate) Set 2:.2 mg/l Ir Set 2:.2 mg/l Ir (Duplicate) No spike: Control No spike: Control (Duplicate) Mean Std. Dev RSD (%) Effect of Chlorite Ion Concentration To investigate the effects of chlorite ion concentration, unstabilized sodium chlorite (Technical Grade, 8%) was added to sodium hypochlorite solutions to achieve a ClO 2 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 ClO 2 and/or at 1 g/l ClO 3 to determine any combined effects of the two anions. Sample solutions prepared in duplicate were incubated at 5 ºC in a water bath. In addition, a similar sample set was incubated at 3 ºC and used to confirm any observable effects of 46

71 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 5 ºC and 3 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. MP = OCl ] [ ClO ] (19) [ 3 (a) OCl g/l Decomposed M MP ( Control ).26 M MP + 15 g/l ClO M MP + 15 g/l ClO g/l ClO M MP + 1 g/l ClO Days ClO 3 g/l Formed (b) OCl g/l Decomposed M MP ( Control ).29 M MP + 15 g/l ClO M MP + 15 g/l ClO 2 +1 g/l ClO M MP + 1 g/l ClO Days ClO 3 g/l Formed 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) 5 ºC, (b) 3 ºC 47

72 As expected, hypochlorite ion solutions at 15 g/l ClO 2 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 2: OCl + ClO 2 ClO 3 + Cl k ClO3 (fast) (2) 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 twofold. 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 smoothedline plots of perchlorate ion formation at 5 and 3 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 2. The observed similarities in changes of the molar product of the chlorate ionspiked sample and chlorite/chlorate ionspiked 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. 48

73 (a) (b) ClO 4 mg/l Formed ClO 4 mg/l Formed M MP ( Control ).26 M MP + 15 g/l ClO M MP + 15 g/l ClO g/l ClO M MP + 1 g/l ClO Days.29 M MP ( Control ).29 M MP + 15 g/l ClO M MP + 15 g/l ClO g/l ClO M MP + 1 g/l ClO Days 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) 5 ºC, (b) 3 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 2. Therefore, chlorite ion is not directly involved in the formation of the perchlorate ion. 49

74 (a) Molar Product M MP + 1 g/l ClO M MP + 15 g/l ClO g/l ClO Days ClO 4 mg/l Formed (b) Figure 2. Molar Product M MP + 1 g/l ClO 3 2.7M MP + 15 g/l ClO g/l ClO Days Plots of perchlorate ion formation in solutions at various initial concentrations of chlorite ion and/or chlorate ion, (a) 3 ºC, (b) 5 ºC 75 ClO 4 mg/l Formed 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 (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ), and bromide ion were investigated. Samples were grouped into the following sets: Set 1: ClO 2 at 15 g/l + Transition Metals Spike at.2 mg/l Set 2: ClO 2 at 15 g/l + ClO 3 at 1 g/l + Transition Metals Spike at.2 mg/l Set 3: ClO 2 at 15 g/l + Br at 15 g/l 5

75 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 5 ºC. The reaction between hypochlorite ion and bromide ion, in the ph 114 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 BrO 2 + Br (22) BrO 2 + OBr BrO 3 + Br (23) Figure 21 shows overlaid smoothedline 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 smoothedline 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, PerlmutterHayman, B. and Stein, G. "The Kinetics of the Decomposition of Alkaline Solutions of HypobromiteSpecific Ionic Effects on Reaction Rate" J. Phys. Chem. 1959, 63,

76 OCl g/l Decomposed No Spike (Control: 69 g/l OCl + 27 g/l ClO 3 ) Set 1: 15 g/l ClO mg/l Me 8 Set 2: 15 g/l ClO g/l ClO mg/l Me Set 3: 15 g/l ClO g/l Br Days Figure 21. Decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions spiked with bromide, chlorite, and transition metal ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) ClO 4 mg/l Formed No Spike (Control: 69 g/l OCl + 27 g/l ClO 3 ) Set 1: 15 g/l ClO mg/l Me Set 2: 15 g/l ClO g/l ClO mg/l Me Set 3: 15 g/l ClO g/l Br Days Figure 22. Formation of perchlorate ion at 5 ºC in solutions spiked with bromide, chlorite, and transition metal ions (Me = Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ) Figure 22 shows that the nonspiked 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 52

77 perchlorate ion than the control solution, due to the higher concentration of chlorate ion than in the control solution. However, because hypochlorite ion was catalyticallydecomposed, 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 2224: OCl + Br OBr + Cl (21) OBr + OBr BrO 2 + Br (22) BrO 2 + OBr BrO 3 + 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: BrO 3 spike at 15 g/l Set 3: Br + BrO 3 spike at 15 g/l Set 4: Br spike at 3 g/l Set 5: BrO 3 spike at 3 g/l Set 6: Br + BrO 3 spike at 3 g/l Figure 23 shows overlaid plots of hypochlorite ion decomposition in various hypochlorite ion solutions at ph 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. 53

78 Figure 23. OCl g/l Decomposed No Spike (Control: 74 g/l OCl + 27 g/l ClO 3 ) Set 1: 15 g/l Br Set 2: 15 g/l BrO 3 Set 3: 15 g/l Br + BrO 3 Set 4: 3 g/l Br Set 5: 3 g/l BrO 3 Set 6: 3 g/l Br + BrO Days Decomposition of hypochlorite ion at 5 ºC in solutions spiked with bromide and bromate ions at ph~12.5 BrO 3 g/l Formed No Spike (BrO 3 at 7 mg/l) Set 1: 15 g/l Br Set 4: 3 g/l Br Set 2: 15 g/l BrO 3 Set 5: 3 g/l BrO 3 Set 3: 15 g/l Br + BrO 3 Set 6: 3 g/l Br + BrO Days Figure 24. Formation of bromate ion at 5 ºC in solutions spiked with bromide 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 54

79 bromate ion. Furthermore, Figure 25 show that perchlorate ion formation in bromate ion spiked samples was identical to control sample (nonspiked 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. ClO 4 mg/l Formed Figure 25. No Spike (Control: 74 g/l OCl + 27 g/l ClO 3 ) Set 1: 15 g/l Br Set 2: 15 g/l BrO 3 Set 3: 15 g/l Br + BrO 3 Set 4: 3 g/l Br Set 5: 3 g/l BrO 3 Set 6: 3 g/l Br + BrO Days Formation of bromate ion at 5 ºC in solutions spiked with bromide and/or bromate ions at ph~ 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. 55

80 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, 1 2 μ = c i zi (24) 2 i where, μ is ionic strength (mol/l), c i concentration (mol/l) of the ith species, and z i is its charge. 86 For illustration purposes, Bleach 21, 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 21: initial concentration of 8 g/l OCl,.2 g/l ClO 3, ph = 12.5, and temperature = 4 ºC. 9 Figure 26. OCl g/l Decomposed g/l OCl g/l Cl 8 g/l OCl + 16 g/l Cl 8 g/l OCl g/l Cl Days Decomposition of hypochlorite ion at 4 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 4 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,

81 ClO 3 g/l Formed g/l OCl g/l Cl 8 g/l OCl + 16 g/l Cl 8 g/l OCl g/l Cl Days Figure 27. Formation of chlorate ion at 4 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. ClO 4 mg/l Formed No Spike (Control:.48 M OCl +.62 M ClO 3 ) Cl Spike at 25 g/l Cl Spike at 5 g/l Cl Spike at 1 g/l Day Figure 28. Formation of perchlorate ion at 4 ºC in solutions at various initial concentrations of chloride ion 57

82 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 1113 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 acidcatalyzed decomposition of hypochlorite ion begins to occur, such that below ph 9 the decomposition of hypochlorite becomes thirdorder 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.9 M OCl solutions were adjusted with hydrochloric acid and incubated at 4 ºC. Figure 29 shows overlaid decomposition plots of the hypochlorite ion and formation plots of the chlorate ion. Figure 3 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 3. At ph 9, the decomposition of hypochlorite ion was too rapid (in a matter of hours 1.4 M OCl decomposed to <.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 58 Region". 58

83 (a) [OCl ] M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph Days [ClO 3 ] (b) [OCl ].9 M OCl +.6 M ClO 3 ph 13.9 M OCl +.6 M ClO 3 ph 11.9 M OCl +.6 M ClO 3 ph Days Figure 29. Decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in (a) 1.4 M OCl, (b).9 M OCl solutions at various initial ph [ClO 3 ] 59

84 (a) ClO 4 mg/l Formed M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph Days Figure 3. (b) ClO 4 mg/l Formed M OCl +.6 M ClO 3 ph 13.9 M OCl +.6 M ClO 3 ph 11.9 M OCl +.6 M ClO 3 ph Days Formation of perchlorate ion at 4 ºC in (a) 1.4 M OCl, (b).9 M OCl solutions at various initial ph As a followup experiment to investigate the effects of ph in 911 region, aliquots of an OSG sodium hypochlorite solution were adjusted to different ph values using sodium hydroxide. Samples solutions at.12 M OCl were adjusted to ph 9.35, 1.65, 11.9, and Because decomposition of.12 M OCl solution is quite slow, samples were incubated at 6 ºC. Figure 31 shows overlaid decomposition plots of hypochlorite ion and formation plots of chlorate ion for these samples. 6

85 (a) (b) ClO 3 g/l Formed OCl g/l Decomposed M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph Days.12 M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph Days Figure 31. Overlaid plots of (a) hypochlorite ion decomposition; (b) chlorate ion formation at 6 ºC.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 standpoint, 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. 61

86 ClO 4 μg/l Formed M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph M OCl +.6 M ClO 3 ph Days Figure 32. Overlaid plots of perchlorate ion formation at 6 ºC.12 M OCl solutions at various initial ph 3.1 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 acidcatalyzed. 62

87 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. OCl + ClO3 ClO4 + Cl (25) Because perchlorate ion formation was dependent on both concentrations of hypochlorite and chlorate ions, the rate law can be expressed by Equation 26. d[ ClO ] Rate ] dt = 4 m p = k2[ OCl ] [ ClO3 (26) By taking the natural log of both sides of Equation 26, an expression shown in Equation 27 is obtained: ln( Rate ) = ln k2 + m ln[ OCl ] + p ln[ ClO3 ] (27) The rate of perchlorate ion formation and the concentrations of hypochlorite and chlorate ions can be measured experimentally. The secondorder rate constant (k 2 ) 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 63

88 the fact that the rate of hypochlorite ion decomposition and the rate of chlorate ion formation are temperaturedependent, 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 1213, where the acidcatalyzed decomposition of hypochlorite ion is at minimum. 9 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[clo 4 ]/dt) vs. ln [ClO 3 ] In this set of experiments, hypochlorite ion solutions were prepared in duplicate with at least 7 g/l OCl, as the initial concentration, while adding chlorate ion at 5 g/l, 1 g/l, and 15 g/l. A control hypochlorite ion solution containing low chlorate ion concentration was used. Samples were incubated at 3, 4, 5, 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 3 ºC. Figures 34 and 35 show the effects of chlorate ion concentration during incubation experiments at 4 ºC and at 5 ºC. Note: the downward sloping smoothed lines represent hypochlorite ion decomposition, and the upwardsloping curves show formation of chlorate ion in Figures 3335 (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 nonspiked hypochlorite ion solutions, and thus the initial concentration of 89 Adam and Gordon "Hypochlorite ion decomposition: Effects of temperature, ionic strength, and chloride ion". 9 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 64

89 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. (a) (b) ClO 4 mg/l 84 g/l OCl + 15 g/l ClO g/l OCl + 63 g/l ClO 3 8 g/l OCl g/l ClO 3 77 g/l OCl g/l ClO Days OCl g/l Decomposed g/l OCl + 15 g/l ClO 3 82 g/l OCl + 63 g/l ClO 3 8 g/l OCl g/l ClO 3 77 g/l OCl g/l ClO Days ClO3 g/l Formed Figure 33. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 3 ºC in solutions at various initial concentrations of chlorate ion at ph ~12.5 ClO 4 mg/l (a) 7 g/l OCl + 28 g/l ClO 3 7 g/l OCl + 51 g/l ClO 3 72 g/l OCl + 15 g/l ClO 3 7 g/l OCl g/l ClO Days OCl g/l Decomposed (b) 7 g/l OCl + 28 g/l ClO 3 7 g/l OCl + 51 g/l ClO 3 72 g/l OCl + 15 g/l ClO 3 7 g/l OCl g/l ClO Days ClO 3 g/l Formed Figure 34. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in solutions at various initial concentrations of chlorate ion at ph ~

90 (a) (b) ClO 4 mg/l g/l OCl + 13 g/l ClO 3 84 g/l OCl + 62 g/l ClO 3 83 g/l OCl + 11 g/l ClO 3 8 g/l OCl g/l ClO Days OCl g/l Decomposed g/l OCl + 13 g/l ClO 3 84 g/l OCl + 62 g/l ClO 3 83 g/l OCl + 11 g/l ClO 3 8 g/l OCl g/l ClO Days ClO 3 g/l Formed Figure 35. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at various initial concentrations of chlorate ion at ph ~12.5 Table 15. Decomposition of hypochlorite ion at 3 ºC in solutions, at various initial concentrations of chlorate ion ([ClO 3 ] ) at ph ~12.5 [OCl ] (mol/l) Days [ClO 3 ] 15 g/l [ClO 3 ] 63 g/l [ClO 3 ] 112 g/l [ClO 3 ] 159 g/l Mean (mol/l) Std. Dev. RSD (%) 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 thirtyfour 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 ] = M and [ClO 3 ] =.178 M increased from 7.1 mg/l ClO 4 to mg/l ClO 4, 66

91 based on the average of duplicate samples. Thus, 53.8 μmol of ClO 4 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 leastsquares lines of natural log of the rate of perchlorate ion formation versus natural log of chlorate ion concentration over different incubation periods at 3 ºC. Figure 36. ln( d[clo 4 ]/dt ) Days 11 y = 1.67x 1.37 R 2 =.999 Days 1121 Days ln[clo 3 ] y =.984x 1.44 R 2 = 1. y = 1.3x 1.35 R 2 =.999 Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 3 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 [ClO 3 ] for chlorate ion spiked experiments at 4 and 5 ºC, respectively. 67

92 ln( d[clo 4 ]/dt ) Day 7 Day 714 y = 1.7x 9.24 R 2 =.998 y = 1.83x 9.41 R 2 = ln[clo 3 ] Figure 37. Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 4 ºC in solutions at constant hypochlorite ion at ph ~12.5 Figure 38. ln( d[clo 4 ]/dt ) Day to 1 y =.981x.354 R 2 =.996 Day 1 to 2 y = 1.44x.568 R 2 = ln[clo 3 ] Fitted natural log lines of the rate of perchlorate ion formation as a function of chlorate ion concentration at 5 ºC in solutions at constant hypochlorite ion at ph ~12.5 The slope of the lines in Figures 3638 represents the reaction order with respect to chlorate ion concentration (p), while the intercept is the sum of ln(k 2 ) and m ln[ocl ] (Equation 28). ln( Rate ) = ln k2 + m ln[ OCl ] + p ln[ ClO3 ] (28) 68

93 The slope and Rsquared values (R 2 ) for the chlorate ion orderfitting at various temperatures are summarized in Table 16. As can be observed from Table 16, fitting ln[rate] versus ln[clo 3 ] produces a linear correlation with an average R 2 =.994, and average slope of 1.35 ±.42 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 R Mean Std. Dev RSD Reaction Order with Respect to Chlorate Ion: ln (d[clo 4 ]/dt) vs. ln [OCl ] In this set of experiments, aliquots of concentrated stock hypochlorite ion solutions, with at least 7 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 3, 4, 5, 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 3 ºC. Figures 4 and 41 show the effects of chlorate ion concentration during incubation experiments at 4 ºC and at 5 ºC. Note: the downwardsloping, smoothed lines in Figures 3941 (b) represent hypochlorite ion decomposition, and the upwardsloping, smoothed lines show formation of chlorate ion. For clarity some plots were omitted in Figure 4 (b). As was 69

94 predicted, the rate of perchlorate ion formation increases with an increase in concentration of hypochlorite ion. (a) (b) ClO 4 mg/l Formed Figure g/L OCl + 63g/L ClO 3 51g/L OCl + 65g/L ClO 3 11g/L OCl + 65g/L ClO Days OCl g/l Decomposed g/L OCl + 63g/L ClO 3 51g/L OCl + 65g/L ClO 3 11g/L OCl + 65g/L ClO Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 3 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~ ClO 3 g/l Formed (a) (b) ClO 4 mg/l Formed 7 g/l OCl + 51 g/l ClO 3 6 g/l OCl + 51 g/l ClO g/l OCl + 52 g/l ClO 3 45 g/l OCl + 51 g/l ClO g/l OCl + 52 g/l ClO g/l OCl + 52 g/l ClO 3 9 g/l OCl + 51 g/l ClO Days g/l OCl Decomposed g/l OCl + 51 g/l ClO 3 52 g/l OCl + 52 g/l ClO 3 45 g/l OCl + 51 g/l ClO 3 25 g/l OCl + 52 g/l ClO Days g/l ClO 3 Formed Figure 4. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 4 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~12.5 7

95 (a) (b) ClO 4 mg/l Formed g/L OCl + 62g/L ClO 3 52g/L OCl + 61g/L ClO 3 1g/L OCl + 56g/L ClO Days g/l OCl Decomposed g/L OCl + 62g/L ClO 3 52g/L OCl + 61g/L ClO 3 1g/L OCl + 56g/L ClO Days g/l ClO 3 Formed Figure 41. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at various initial concentrations of hypochlorite ion at ph ~12.5 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 3 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 3 ºC in solutions at constant chlorate ion concentration at ph ~12.5 Days [OCl ] M [OCl ].982 M [OCl ].26 M Mean Std. Dev. RSD

96 Figures 4245 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.. Figure 42. ln( d[clo 4 ]/dt ) Days 11 y = 1.4x 11.3 R 2 = Days 2134 y = 1.62x 11.2 R 2 = ln[ocl ] Days 1121 y = 1.87x 11. R 2 =.996 Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 3 ºC in solutions at constant chlorate ion at ph ~12.5 ln( d[clo 4 ]/dt ) Day 7 Day 714 Day 1421 y = 1.14x 7.95 R 2 =.979 y = 1.72x 9.99 R 2 =.985 y = 1.67x 9.87 R 2 = ln[ocl ] Figure 43. Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 4 ºC in solutions constant in chlorate ion, with ph ~

97 . 1. Day to 1 y = 1.45x 1.54 R 2 =.991 Figure 44. ln( d[clo 4 ]/dt ) Day 1 to 2 Day 2 to 3 y = 1.18x 1.32 R 2 =. 986 y = 1.25x 2.1 R 2 = ln[ocl ] Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 5 ºC in solutions at constant chlorate ion at ph ~12.5 ln( d[clo 4 ]/dt ) Day to 1 y = 1.27x 6.13 R 2 = ln[ocl ] Figure 45. Fitted natural log lines of the rate of perchlorate ion formation as a function of hypochlorite ion concentration at 75 ºC in solutions at constant chlorate ion at ph ~12.5 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 ] [ClO 3 ] = molar product) were conducted. In this set of experiments, solutions 73

98 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 3 and 5 ºC. Note: the downward sloping smoothed lines and the upwardsloping curves in Figures 4647 (b) represent hypochlorite ion decomposition and chlorate ion formation, respectively. (a) (b) ClO 4 mg/l Formed M OCl +.18 M ClO 3.99 M OCl +.3 M ClO 3.2 M OCl M ClO Days Figure 46. [OCl ] M OCl +.18 M ClO 3.99 M OCl +.3 M ClO 3.2 M OCl M ClO Days Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 3 ºC in solutions at constant molar product at ph ~ [ClO 3 ] 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 secondorder, but rather a higher order process due to larger dependence on concentration of hypochlorite ion, or there is an experimental variable not accounted for. 74

99 (a) (b) ClO 4 mg/l Formed M OCl +.16 M ClO M OCl +.26 M ClO 3.2 M OCl M ClO Days [OCl ] M OCl +.16 M ClO M OCl +.26 M ClO 3.2 M OCl M ClO Days [ClO 3 ] Figure 47. Overlaid plots of (a) perchlorate ion formation; (b) decomposition of hypochlorite ion and formation of chlorate ion at 5 ºC in solutions at constant molar product at ph ~ 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 secondorder 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 tworeaction 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 firstorder and secondorder in hypochlorite ion. Where k 1 is a secondorder rate constant, and k 2 is a thirdorder rate constant. d[ ClO4 ] 2 Rate = = k1[ OCl ][ ClO3 ] + k2[ OCl ] [ ClO3 ] (28) dt A consecutive reaction pathway, given by Equation 29, may involve a competitive preequilibrium reaction with another hypochlorite ion to form an 75

100 intermediate. This preequilibrium reaction is followed by decomposition of the intermediate to form either reactants or products. Where k a, is a rate constant that drives the reaction to form products, and the value of rate constant, k b, determines the extent of the forward reaction. 2 d[ ClO4 ] k [ OCl ] [ ClO3 ] Rate = = a (29) dt 1+ kb[ OCl ] 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 k 1 and k 2. 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 3 is obtained. Rate = k1 + k2[ OCl ] (3) [ OCl ][ ClO3 ] When the [Rate / ([OCl ] [ClO 3 ])] is plotted versus [OCl ] (for constant chlorate ion experiments), a linear correlation will provide values of k 2 (the slope of the line) and k 1 (the intercept). Figure 48 shows linear plots at different temperatures. Table 18 shows the values of k 2 and k 1 at each temperature, calculated from the slope and intercept of the fitted lines in Figure

101 (a) 3 ºC Rate/[OCl ][ClO 3 ] y = x R 2 = [OCl ] (b) 4 ºC (c) 5 ºC Rate/[OCl ][ClO 3 ] y = x R 2 = [OCl ] Figure 48. Rate/[OCl ][ClO 3 ] y = x R 2 = [OCl ] Parallel reaction pathway linear fitted plots of (a) 3 ºC experiment; (b) 4 ºC experiment; (c) 5 ºC experiment; in solutions at constant chlorate ion and various inial voncentrations of hypochlorite ion at ph ~

102 Table 18. Parallel reaction pathway experimental rate constants in solutions, at various hypochlorite ion and constant chlorate ion, at ph ~12.5 T, ºC k M 2 d 1 k M 1 d 1 R 2 k 2 /k An initial analysis of k 2 and k 1 values shows that both rate constants increase with temperature; however, the ratio of k 2 /k 1 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 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 k a and k b. Equation 29 can be rearranged by dividing both sides of the equation by the concentration of hypochlorite ion to obtain Equation 31: 2 Rate k [ ] [ ] = a OCl ClO3 (31) [ OCl ] (1 + kb[ OCl ])[ OCl ] If, Equation 31 is inverted and both sides are multiplied by the concentration of chlorate ion, Equation 32 results: [ OCl ][ ClO3 ] (1 + k [ ])[ ][ 3 ] = b OCl OCl ClO (32) Rate 2 ka[ OCl ] [ ClO3 ] Equation 33 results from cancelling the common terms on the right side of Equation 32. [ OCl ][ ClO3 ] 1 k = + b (33) Rate k k a[ OCl ] a By plotting the quantity [OCl ] [ClO 3 ]/(Rate) versus 1/[OCl ], the values of k a and k b can be determined from the leastsquares lines: the slope equals 1/k a and the 78

103 intercept is k b /k a (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 k a and k b at each temperature, calculated from the slope and intercept of the fitted lines in Figure 49. (a) 3 ºC (b) 4 ºC [OCl ][ClO 3 ] / Rate [OCl ][ClO 3 ] / Rate y = x R 2 = /[OCl ] y = x R 2 = /[OCl ] (c) 5 ºC [OCl ][ClO 3 ] / Rate y = x R 2 = /[OCl ] Figure 49. Consecutive reaction pathway linear fitted plots of (a) 3 ºC experiment; (b) 4 ºC experiment; (c) 5 ºC experiment; in solutions at constant chlorate ion and various hypochlorite ion at ph ~

104 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 k a 1 6 M 2 d 1 k b M 1 d 1 R 2 k b / k a , , ,7 The linear fit of the experimental data to a consecutive reaction pathway for the 3 ºC experiment, shown in Figure 49, has a low Rsquared value of.951. In addition, as shown in Table 19, the secondorder rate constant, k b, is highlyvariable 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 2, it was discovered that the differences in ionic strength of these solutions were not significant. Thus, the experimentallymeasured reaction order with respect to chlorate ion of one must be valid. 8

105 Table 2. Ionic strength (μ) of hypochlorite ion solutions at various chlorate ion at 4 ºC experiments (TDS = Total Dissolved Solids) [ClO 3 ], mol/l [OCl ], mol/l μ, mol/l ph TDS, g/l Mean Std. Dev RSD Incubation experiments at 4 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 ( mol/l). Table 21. Ionic strength (μ) of hypochlorite ion solutions at various hypochlorite ion at 4 ºC experiments (TDS=Total Dissolved Solids) [OCl ], mol/l [ClO 3 ], mol/l μ, mol/l ph TDS, g/l Mean Std. Dev RSD 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 5. 81

106 (a) y = x x R 2 =.999 ln( d[clo 4 ]/dt ) Ionic Strength (μ), mol/l (b) y = x x R 2 =.997 ln( d[clo 4 ]/dt ) [OCl ], mol/l Figure 5. Rate of perchlorate ion formation as a function of (a) ionic strength; (b) concentration of hypochlorite ion, at 4 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 5. At low ionic strength and low concentrations of hypochlorite ion, the rate of perchlorate ion formation changes more slowly. At ionic strength above 2. M and.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. 82

107 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, secondorder reaction that takes place between hypochlorite and chlorate ions. The rate of perchlorate ion formation is shown in Equation 27: d[ ClO k OCl m ClO p Rate 4 ] = = 2[ ] [ 3 ] (27) dt The order with respect to chlorate ion has been determined to be firstorder, 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 firstorder. Thus, a simpler model based on reaction that is firstorder in both (m, p = 1) chlorate and hypochlorite ions is obtained, allowing determination of the secondorder rate constant by Equation 34. Rate k 2 = (34) [OCl ] [ClO3 ] The fitting of the experimental rate of perchlorate ion formation as a function of ionic strength from the 4 ºC experiments, as shown in Figure 5, demonstrated linearity at ionic strengths greater than 2 M. However, before any further conclusions can be made, an expression relating the secondorder rate constant, k 2, and ionic strength is needed. 83

108 4.4.1 Dependence of the SecondOrder Rate Constant on the Ionic Strength Transition state theory (developed by Henry Eyring 91 ) 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. k OCl + ClO 2 k 3 [OCl ClO3 ] ClO4 + Cl (35) Where, k 2, is the experimental secondorder rate constant, and [OCl ClO 3 ] denotes the activated complex. The formation of the quasi equilibrium between reactants and [OCl ClO 3 ] intermediate can be described by an equilibrium constant K, shown in Equation 36: [OCl ClO = 3 ] γ K OCl [OCl ] [ClO3 ] γ γ ClO3 where, γ is the activity coefficient. Thus the rate of this reaction can be defined as shown in Equation 37. d[ ClO4 ] Rate = = k2[ OCl ] [ ClO3 ] = k [OCl ClO3 ] (37) dt By defining the quantity [OCl ClO 3 ] using Equation 36 and substituting into Equation 37, the secondorder rate constant, k 2, can be expressed by Equation 38. (36) OCl ClO3 γ γ k 2 = k K (38) where the quantity k K, is defined as the rate constant at infinite dilution, and is commonly defined as k ref or k o. Thus Equation 38 really is the same as the Brønsted Bjerrum Equation 92, shown by Equation 39: γ OCl ClO3 γ γ k 2 = k o (39) γ Thus, by taking the log of both sides of the Equation 39, Equation 4 results. 91 Espenson, J. H. Chemical Kinetics and Reaction Mechanisms, 2nd ed.; McGrawHill, Inc., Espenson Chemical Kinetics and Reaction Mechanisms. 84

109 OCl ClO3 γ γ log( k 2) = log( k o ) + log( ) (4) γ A reduced form of a DebyeHückel Equation by Güntelbuerg 93, shown in Equation 41, can be used for all electrolytes at 25 ºC to relate the activity coefficient to the ionic strength (γ i denotes activity coefficient for each species i). A Z log( γ i ) = 1+ 2 i μ (41) μ where Z 1 and Z 2 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.88m, thus the modified Equation 41 by Gugenheim, which adds a linear term 94 for solutions with ionic strength above.1 M, by Equation 42, is needed. 2 i μ i + A Z log( γ ) = 1+ where b is an adjustable parameter 95, Davies in his modification of Equation 41 used the value of.2 for b and.5 for the term A at 25 ºC. 96 By constructing three equations like Equation 42 and combining them into Equation 4, Equation 43 results: 2AZ Z OCl ClO μ 3 log( k2) = log( ko ) + ( bμ) (43) 1+ μ 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. μ bμ (42) 93 Robinson, R. A. and Stokes, R. H. Electrolyte Solutions, 2nd ed.; Butterworths Publications Limited: London, 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, 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,

110 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 ionion pairing becomes increasingly dependent on ionic strength above.1 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 ionion interactions. 98 Equation 43 incorporates both the effect based on increased activity of ionic species, described by the term μ /( 1+ μ) and the effect of ionion interactions, described by the termb μ. Both terms were plotted as a function of ionic strength, shown in Figure 51, so that their relative contributions can be evaluated. Relative Contribution bμ μ /( 1+ μ) μ, mol/l Figure 51. Plot of the μ /( 1+ μ) and bμ terms as a function of ionic strength. Note: value of.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. 86

111 As can be observed from Figure 51, the contribution of the ionion interaction term 99 becomes more significant at μ 1 M, by assuming b =.5. Thus, an empirical relationship between ionic strength, the secondorder rate constant, k 2, and the rate constant at infinite dilution, k, is approximated by Equation 44: log(k2 ) = b μ + log(ko ) (44) The values of k 2 were calculated for experiments conducted at different temperatures by using Equation 44 and plotted against the ionic strength of solutions, shown in Figure ºC: y =.738x 1.1 R 2 = 1. 4 ºC: y =.838x 9.66 R 2 = ºC: y =.788x 9. R 2 =.978 log(k 2 ) Figure μ, mol/l Overlaid linear plots of log of secondorder rate constant versus ionic strength in solutions with various initial concentrations of hypochlorite ion at different temperatures The slopes of the leastsquares lines are equal to the b term, and the intercepts are log of the secondorder rate constant at infinite dilution. The log (k o ) 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 secondorder rate constant of perchlorate ion formation. Thus, Equation Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 87

112 provides the simplified approximation of the changes in the secondorder rate constant as a function of the ionic strength at different temperatures. Table 22. Slopes and intercepts of leastsquares lines shown in Figure 52 Temperature (ºC) Slope Intercept (log k ) k o in M 1 s 1 (x 1 12 ) Mean.788 Std. Dev..5 RSD Dependence of the SecondOrder 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 secondorder rate constant at infinite dilution. Figure 53. ClO 4 mg/l Formed ºC [OCl ] = 84.2 g/l, [ClO 3 ] = 14.8 g/l 4 ºC [OCl ] = 7.5 g/l, [ClO 3 ] = 28.1 g/l 5 ºC [OCl ] = 85.3 g/l, [ClO 3 ] = 13. g/l 75 ºC [OCl ] = 75. g/l, [ClO 3 ] = 24.2 g/l Days Smoothline plots of the perchlorate ion formation as a function of time in solutions at similar initial hypochlorite and chlorate ions and various temperatures 88

113 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 equation 1 is shown in Equation 47 below: ΔS ΔH k ( ) ( ) k b R RT o = T e e h (45) k o = Calculated using Equation 27 M 1 s 1 k b = Boltzmann constant 1.381x1 23 J K 1 h = Planck constant 6.626x1 34 J s R = Real Gas constant 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: k k o b ΔS ΔH ln( ) = ln( ) + (46) T h R RT To determine ΔH and ΔS, ln(k o /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 (k b / h) + (ΔS / R)]. Temperatures in Table 2 were converted from ºC to Kelvin to determine values of ln(k o /T) and 1/T terms. Figure 54 shows a plot of experimentally determined ln(k o /T) vs 1/T. 1 Espenson Chemical Kinetics and Reaction Mechanisms. 89

114 Figure 54. ln (k o /T) y = 12165x R 2 = /T, K Linear plot of ln(k o /T) as a function of (1/T) Based on three points the R 2 of the line is.983. The relatively low rsquared correlation coefficient of.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 2, the R 2 correlation coefficient of.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 = 11 kj/mol and ΔS = 16 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. 11 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 secondorder rate constant at infinite dilution and thermodynamic activation parameters, shown in Equation x k = T e RT e R o (47) 11 Espenson Chemical Kinetics and Reaction Mechanisms. 9

115 4.4.3 Combining the Effects of the Ionic Strength and Temperature on the SecondOrder 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 firstorder concentrations of hypochlorite and chlorate ions, ionic strength, and temperature. The secondorder 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.788 and Equation 47 into Equation 44, the combined effects of ionic strength and temperature are given by Equation 48: 5 1.1x log(k ).788( ) log( T e RT e R 2 = μ + ) (48) The secondorder rate constant of perchlorate ion formation for solutions of any ionic strength in the range of mol/l and temperature range 35 ºC can be calculated by using Equation 48. Table 23 shows the experimentally observed values of secondorder rate constant, k 2, and predicted values by using Equation 48. Table 23. Experimental and predicted secondorder rate constants at variable ionic strength and temperature (k exp = experimental k 2 ; k pred = predicted k 2 ) T μ k exp k pred % Average R 2 (ºC) (M) (M 1 d ) (M 1 d ) Error % Error log(k 2 ) vs. μ

116 The difference between the experimentallyobserved and predicted value of k 2 is reported as percent error in Table 23. Given the relatively low R 2 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 secondorder rate constant with a strong dependence on the ionic strength. The dependence of the secondorder 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 secondorder rate constant is provided to utilities to predict how to decrease the amount of perchlorate ion formed during storage of hypochlorite ion solutions. 92

117 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 secondorder rate constant, ionic strength, and temperature (Equation 5) with the use of hypochlorite ion decomposition Bleach 21 to predict perchlorate ion formation in bulk sodium hypochlorite solutions. The objective for the predictive model is to have an average error of ± 1 %. 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 21, 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 21, multiple commercial hypochlorite ion solutions were tested, and an average error of ± 5 % was reported between experimentally measured decomposition and predicted decomposition. 12 In the design of the experiments conducted in the current work, Bleach 21 was used to predict decomposition of hypochlorite ion and formation of chlorate ion. Figure 55 shows the decomposition of hypochlorite ion (downwardsloping curves) and formation of chlorate ion (upwardsloping curves) measured experimentally and predicted by Bleach 21 at 3º, 4º, and 5 ºC. 12 Adam An Investigation of Factors Involved In The Decomposition of Sodium Hypochlorite. 93

118 (a) 3 ºC (b) 4 ºC OCl g/l Decomposed Experiment Bleach Days ClO 3 g/l Formed (c) 5 ºC OCl g/l Decomposed Bleach 21 Experiment Days ClO 3 g/l Formed Figure 55. OCl g/l Decomposed Experiment Bleach Days Smoothedline plots of hypochlorite ion decomposition and chlorate ion formation determined experimentally in conjunction with Bleach 21 (Error bars set at ± 1 %), for solutions at (a) [OCl ] = 82 g/l, [ClO 3 ] = 63 g/l at 3 ºC; (b) [OCl ] = 7 g/l, [ClO 3 ] = 51 g/l at 4 ºC; (c) [OCl ] = 83 g/l, [ClO 3 ] = 5 g/l at 5 ºC ClO 3 g/l Formed 94

119 The average % difference between measured hypochlorite ion and predicted values by Bleach 21, at 3 ºC over a 2day period, was 8.8 % ± 4.2 (n = 7); at 4 ºC over 34day period, the average % difference was 2.2 % ± 8.8 (n = 4); and at 5 ºC, the average % difference over 1 days was 6.6 % ± 4.6 (n = 6). Thus, the predictive expression for the secondorder rate constant (Equation 5) was used with Bleach 21, 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 secondorder rate constant. Based on the ionic strength of the hypochlorite ion solutions and the incubation temperature, secondorder rate constants were generated, using Equation x log(k ).788( ) log( T e RT e R 2 = μ + ) (48) The Bleach 21 model was used to predict the decomposition of hypochlorite ion and formation of chlorate ion at various temperatures. The predicted values of k 2, 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. [ 4 d ClO ] Rate = = k2[ OCl ] [ ClO3 ] (27) dt As an example, to simulate the perchlorate ion formation at 7.4 g/l OCl, 5.1 g/l ClO 3, and μ = 6.76 mol/l during a 4 ºC incubation, the value of k 2 was calculated by entering temperature (converted from 4 ºC to Kelvin) and ionic strength into Equation 5. The obtained value of k 2, L mol 1 s 1, was converted to units of L mol 1 d 1, producing a value of The predicted value of k 2, and predicted changes in concentration of hypochlorite and chlorate ions by Bleach 21 were 95

120 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[clo 4 ]/dt as a function of time at 4 ºC Day [OCl ], (mol/l) [ClO 3 ], (mol/l) d[clo 4 ]/dt, (mol/l/day 1 6 ) d[clo 4 ]/dt, (mg/l / day) mg/l ClO 4 produced The summation of incremental increases in perchlorate ion concentration over initial concentration as a function of time is recorded at each dayinterval. Figure 56 shows overlaid plots of predicted perchlorate ion concentration as a function of time and experimentally measured concentration. The smoothedlines represent predicted changes in perchlorate ion formation, with error bars set at a fixed ± 1 %. 96

121 (a) 3 ºC 35 3 [OCl ] = 83 g/l; [ClO 3 ] =5 g/l [OCl ] = 52 g/l; [ClO 3 ] =5 g/l [OCl ] = 1 g/l; [ClO 3 ] =5 g/l ClO 4 mg/l (b) 4 ºC Days 35 3 [OCl ] = 7 g/l; [ClO 3 ] =5 g/l [OCl ] = 6 g/l; [ClO 3 ] =5 g/l [OCl ] = 33 g/l; [ClO 3 ] =5 g/l ClO 4 mg/l Days (c) 5 ºC 35 3 [OCl ] = 83 g/l; [ClO 3 ] =5 g/l [OCl ] = 52 g/l; [ClO 3 ] =5 g/l [OCl ] = 1 g/l; [ClO 3 ] =5 g/l ClO 4 mg/l Days Figure 56. Overlaid smoothedline plots of predicted (Error bars set at ± 1 %) perchlorate ion formation and determined experimentally, for solutions incubated at (a) 3 ºC; (b) 4 ºC; (c) 5 ºC 97

122 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 ± 1 %, 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 5 ºC after several days. To test the model further, several hypochlorite ion solutions were acquired from utilities in United States and incubated at 5 ºC. These results are discussed in the next section. 5.2 Predicted Perchlorate Ion Formation in RealWorld 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 nonquenched sample that was cooled to 4 ºC. The samples were analyzed for bromate, and perchlorate ions by the LCMS/MS method (Chpater 2). ICPMS analysis was used to screen for transition metals ions (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ). 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 ClO 4 (mg/l) BrO 3 (mg/l) Mn (mg/l) Fe (mg/l) Co (mg/l) Ni (mg/l) Cu (mg/l) < < <.1 < 5 <.1 <.1 < < <.1 <.1 < <.1 <.5 <.1.11 < < <.1 <.1 <.1 98

123 Table 26. Chlorate and hypochlorite ions, ph, TDS, and ionic strength in bulk utility hypochlorite ion solutions Utility Solution OCl (g/l) ClO 3 (g/l) ph μ (mol/l) TDS (g/l) Hypochlorite ion solutions were aged at 5 ºC to determine the rate of perchlorate ion formation and decomposition of hypochlorite ion during a 3day incubation study. Overlaid smoothedline plots of perchlorate and chlorate ion formation and hypochlorite ion decomposition in sodium hypochlorite from different utilities are shown in Figure 57. OCl g/l (a) Hypochlorite Ion Decomposition Utility 1 Utility 2 Utility 3 Utility 4 Utility Days ClO 3 g/l (b) Chlorate Ion Formation Utility 1 Utility 2 Utility 3 Utility 4 Utility Days (c) Perchlorate Ion Formation Figure 57. mg/l ClO 4 35 Utility 1 3 Utility 2 Utility 3 25 Utility 4 2 Utility Days Overlaid smoothedline plots of (a) hypochlorite ion decomposition; (b) chlorate ion formation; (c) perchlorate ion formation determined experimentally during incubation at 5 ºC 99

124 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 = 5 ºC from Table 26 were entered into Bleach 21, 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. OCl g/l (a) Utility 1 (Experiment) Bleach Days ClO 3 g/l OCl g/l (b) Utility 2 (Experiment) Bleach Days 7 ClO 3 g/l (c) (d) OCl g/l 12 Utility 3 (Experiment) 5 1 Bleach Figure 58. Days ClO 3 g/l OCl g/l Utility 5 (Experiment) Bleach Days Overlaid smoothedline plots of hypochlorite ion decomposition and chlorate ion formation determined experimentally in conjunction with Bleach 21 (Error bars set at ± 1%), for solutions with (a) [OCl ] = 63 g/l, [ClO 3 ] = 23 g/l;(b) [OCl ] = 111 g/l, [ClO 3 ] = 8.7 g/l; (c) [OCl ] = 89 g/l, [ClO 3 ] = 4.4 g/l; (d) [OCl ] = 97 g/l, [ClO 3 ] = 12 g/l; incubated at 5 ºC ClO 3 g/l 1

125 Figure 58 shows the observed rates of hypochlorite ion decomposition, chlorate ion formation and predicted rates of different utilities agree well within ± 1 % 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 21 predicted concentrations of hypochlorite and chlorate ions as a function of time and the predicted secondorder 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. ClO 4 mg/l (a) Utility 1 (Experiment) Model Days ClO 4 mg/l (b) Utility 2 (Experiment) Model Days ClO 4 mg/l (c) Utility 3 (Experiment) Model Days ClO 4 mg/l Utility 5 (Experiment) Model Days Figure 59. Overlaid smoothedline plots of predicted (Error bars set at ± 1 %) perchlorate ion formation and determined experimentally at 5 ºC in solutions with (a) [OCl ] = 63 g/l, [ClO 3 ] = 23 g/l; (b) [OCl ] = 111 g/l, [ClO 3 ] = 8.7 g/l(c) [OCl ] = 89 g/l, [ClO 3 ] = 4.4 g/l; (d) [OCl ] = 97 g/l, [ClO 3 ] = 12 g/l (d) 11

126 Figure 59 illustrates that the developed perchlorate ion model simulates the perchlorate ion formation in bulk sodium hypochlorite solutions with ph in range of within ± 1 %. 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 21 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 secondorder rate constant for perchlorate ion formation and the rate of hypochlorite ion decomposition. Figure 6 shows overlaid smoothedline plots of the rate of perchlorate ion formation as a function of dilution factor or temperature. As expected, Figure 6 demonstrates that the rate of perchlorate ion formation is reduced markedly when hypochlorite ion solution is either 12

127 diluted or cooled. By reducing the temperature of sodium hypochlorite solution from 35 ºC to 3 ºC, the rate of perchlorate ion formation decreases by a factor of This factor changes slightly as a function of temperature. For example, cooling hypochlorite ion solutions from 5 ºC to 4 ºC reduces the rate by factor of 3.4; however, cooling a 3 ºC solution to 2 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 4 ºC. (a) % Rate Cooling by 5 ºC reduces the rate to 51.4% (by 1.95) Temperature, ºC Figure 6. (b) % Rate At 1:2 dilution, the rate is reduced to 12.7% (by 7.9) Dilution Factor Smoothedline 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,.34 M ClO 3, and μ = 7.5 M; rate at 35 ºC = 1% 1. 13

128 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 ] = 2.53 M, [ClO 3 ] =.34 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 6). 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 costefficient. Even if storing at elevated temperatures is necessary, dilution is a very effective strategy to minimize perchlorate ion formation. Figure 61 (a) shows overlaid, smoothline 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 2.3 M OCl (equivalent to 13 % FAC) and at 1.2 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 watertreatment, 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 ClO 4 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 ClO 4 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.726 mg OCl ). 14

129 (a) [OCl ], M M OCl 1.2 M OCl Days mg/l ClO 4 Figure 61. (b) μg ClO 4 / mg FAC M OCl 1.2 M OCl MA MCL = 2 μg/l ClO Days Overlaid smoothedline plots of (a) predicted decomposition of hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO 4 per mg OCl as a function of time in solutions at 2.3 M OCl and 1.2 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. 13 Thus, in this case, after 75 days of storage at 35 ºC, the concentration of perchlorate ion in treated 13 MassDEP, "Perchlorate in Public Drinking Water", 26, Retrieved 7/21/9, from: 15

130 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. (a) [OCl ], M M OCl 1.2 M OCl Days mg/l ClO 4 Figure 62. (b) μg ClO 4 / mg OCl M OCl 1.2 M OCl Days Overlaid smoothedline plots of (a) predicted decomposition of hypochlorite ion and formation of perchlorate ion; (b) plot of μg ClO 4 per mg OCl as a function of time in solutions at 2.3 M OCl and 1.2 M OCl at 25 ºC 16

131 In this section of the work it has been demonstrated that diluting hypochlorite ion solutions increases the shelflife of the disinfectant solution, while at the same time significantly minimizes formation of perchlorate ion. Cooling diluted hypochlorite ion solutions by as little as 51 ºC, further reduces the formation of perchlorate ion and also increases the shelflife of the disinfectant. Thus, by diluting the bulk hypochlorite ion solutions and storing them below 3 ºC, the formation of perchlorate ion is significantly minimized. 5.4 Application of the Perchlorate Model to OSG Sodium Hypochlorite Solutions Because the onsite 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 onsite sodium hypochlorite generators manufacturers with different capacities, ranging from 1 2, pounds per day. The concentration of bromate, chlorate, and perchlorate ions was determined by the LCMS/MS method. ICPMS analysis was used to determine the concentrations of transition metals ions (Co 2+, Cu 2+, Fe 3+, Mn 2+, and Ni 2+ ). 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

132 Table 27. Transition metals, bromate, chlorate, and perchlorate ions in OSG hypochlorite ion solutions OSG Solution g/l ClO 3 (g/l) BrO 3 (mg/l) ClO 4 (μg/l) Mn (μg/l) Fe (μg/l) Co (μg/l) Ni (μg/l) Cu (μg/l) < 25 < 125 < 25 < 25 < < < 25 < < < 25 < 25 < < < 25 < 25 < < < 25 < 25 < < 25 < 125 < 25 < 25 < < 25 < 125 < 25 < 25 < < 25 < 125 < 25 < 25 < < 5 < 5. < < 5 < 5. < 5. < < 1 < 125 < 1 < 1 < < 1 < 25 < 1 < 1 < 1 Note: Detection limits vary depending on the dilution factor used for analysis Figure 63 shows overlaid smoothedline 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 91. In this ph range the kinetics of acidcatalyzed decomposition of hypochlorite ion becomes faster and dependent on more variables that are outside the scope of Bleach 21. However, the rate of hypochlorite ion decomposition is strongly dependent on the initial concentration of hypochlorite ion and even when incubated at 5 ºC, the rate of decomposition is relatively slow. Overlaid smoothedline plots of perchlorate ion formation are shown in Figure

133 (a) OCl g/l OSG 1 OSG 2 OSG 3 OSG 4 OSG 5 OSG Days Figure 63. (b) OCl g/l Days OSG 7 OSG 8 OSG 9 OSG 1 OSG 11 OSG 12 Overlaid smoothedline plots of hypochlorite ion decomposition (a) OSG solutions 16; (b) OSG solutions 712, 5 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

134 (a) ClO 4 μg/l OSG 1 OSG 2 OSG 3 OSG 4 OSG 5 OSG Days ClO 4 μg/l (OSG 4) Figure 64. (b) ClO 4 μg/l OSG 7 OSG 8 OSG 9 OSG 1 OSG 11 OSG Days Overlaid smoothedline plots of perchlorate ion formation (a) OSG solutions 16; (b) OSG solutions 712, 5 ºC ClO 4 μg/l (OSG 1, 12) 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 secondorder rate constant, using Equation 36. By using the measured values of ionic strength and temperature, the predictive model was used to determine the secondorder rate constants. Table 29 shows the experimental and predicted values of the secondorder rate constants of perchlorate ion formation. 11

135 Table 29. Secondorder rate constants of perchlorate ion formation in OSG hypochlorite ion solutions, experiment vs. model (k 2obs = experimental k 2, k 2cal = predicted k 2, in units of L mol 1 day 1 ) OSG k 2obs k 2cal k 2obs /k 2cal 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 experimentallyobserved values of the secondorder 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. 111

136 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. 14 It also has been reported that perchlorate ion is not a common contaminant in calcium hypochlorite. 15 A sample of calcium hypochlorite, 68 % 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 6 ºC and four weeks at 5 ºC. Figure 65 shows smoothedline 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 LCMS/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 5 ºC and 6 ºC and is shown in Figure Su, Y. S., Morrison, D. T. and Ogle, R. A. "Chemical kinetics of calcium hypochlorite decomposition in aqueous solutions" J. Chem. Health Safety 29, 16, Greiner, Mclellan, Bennet and Ewing "Occurrence of perchlorate in sodium hypochlorite". 112

137 OCl g/l Figure 65. (a) 5 ºC (b) 6 ºC [OCl ] = 44 g/l; [ClO 3 ] =.8 g/l [OCl ] = 23 g/l; [ClO 3 ] =.4 g/l Days ClO 3 g/l OCl g/l [OCl ] = 45 g/l; [ClO 3 ] =.8 g/l [OCl ] = 23 g/l; [ClO 3 ] =.4 g/l Days Smoothedline plots of hypochlorite ion decomposition and chlorate ion formation in calcium hypochlorite solutions (a) incubated at 5 ºC; (b) incubated at 6 ºC ClO 3 g/l (a) 5 ºC (b) 6 ºC ClO 4 mg/l [OCl ] = 44 g/l; [ClO 3 ] =.8 g/l 1 [OCl ] = 23 g/l; [ClO 3 ] =.4 g/l Days Figure 66. ClO 4 mg/l [OCl ] = 45 g/l; [ClO 3 ] =.8 g/l [OCl ] = 23 g/l; [ClO 3 ] =.4 g/l Days Smoothedline plots of perchlorate ion formation in calcium hypochlorite solutions (a) incubated at 5 ºC; (b) incubated at 6 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 6 ºC was measured at.76 μmol/l per day, and at 5 ºC, the rate was.2 μmol/l per day in solution at 23 g/l OCl and.4 g/l ClO 3. Thus, if the temperature is increased by 1 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 3 shows the calculated and measured secondorder rate constants for different calcium hypochlorite solutions. 113

138 Table 3. Secondorder rate constants of perchlorate ion formation in calcium hypochlorite solutions, experiment vs. model (k 2obs = experiment k 2, k 2cal = k 2 predicted in units of L mol 1 day 1 ) T, ºC [OCl ] (mol/l) [ClO 3 ] (mol/l) k obs k cal k 2obs / k 2cal ph Ionic strength The results shown in Table 27 indicate that the observed secondorder rate constants are higher than predicted (as noted in k 2obs / k 2cal 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 hypochlorite 16 (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 Robinson and Stokes Electrolyte Solutions. 114

139 Table 31. Residence time of the sampled distribution waters Utility Res. Time (A) (hrs) Res. Time (B) (hrs) Table 32. Perchlorate ion in raw, finished, and distribution waters ClO 4 Raw Finished Time A Time B Utility (μg/l) (μg/l) (μg/l) (μg/l) 1 < < <.5 <.5 <.5 <.5 3 <.5 <.5 <.5 <.5 4 <.5 <.5 <.5 < Table 33. Bromate ion in raw, finished, and distribution waters BrO 3 Raw Finished Time A Time B Utility (μg/l) (μg/l) (μg/l) (μg/l) 1 < <.5 <.5 <.5 < < < Table 34. Chlorate ion in raw, finished, and distribution waters ClO 3 Raw Finished Time A Time B Utility (μg/l) (μg/l) (μg/l) (μg/l) < 3. < 3. 4 < < 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 115

140 concentration above the USEPA established interim health advisory level 17 of 15 μg/l. Bromate ion concentration was measured below an MCL of 1 μg/l in all water samples (Table 33). However, chlorate ion was present in several samples above the guideline MCL of 7 µg/l (Table 34), set by the World Health Organization. 18 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 ClO 3, which indicates that this solution has been decomposing. Solution from Utility 3 contained 89. g/l OCl and 4.37 g/l ClO 3, 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 17 USEPA "Interim Drinking Water Health Advisory for Perchlorate" 18 WHO Guidelines for Drinking Water Quality; World Health Organization: Singapore, 26, 116

141 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 ng ClO 4 per mg FAC ng BrO 3 per mg FAC μg ClO 3 per mg FAC FAC (g/l ) OSG Sodium Hypochlorite ng ClO 4 per mg FAC ng BrO 3 per mg FAC μg ClO 3 per mg FAC FAC (g/l ) Calcium Hypochlorite ng ClO 4 per mg FAC ng BrO 3 per mg FAC μg ClO 3 per mg FAC FAC (g/l ) As an alternative, small drinkingwater utilities may consider the use of calcium hypochlorite for disinfection treatments. Calcium hypochlorite contained the lowest ng ClO 4 per mg FAC, and one of lowest ng BrO 3 per mg FAC and μg ClO 3 per mg FAC, as can be seen in Table

142 5.7 Conclusions The use of Equation 5 to determine the value of secondorder rate constant as a function of the ionic strength of solution at specific temperatures has been validated. A difference of less than ± 1 % between the experimentallymeasured decomposition of bulk hypochlorite ion and that predicted by Bleach 21 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: 5 1.1x log(k ).788( ) log( T e RT e R 2 = 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. μg/l). However, the results revealed that the levels of chlorate ion are considerably higher than the WHO MCL 19 of 7 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. 19 WHO Guidelines for Drinking Water Quality 118

143 CHAPTER 6. CONCLUSIONS 6.1 Summary A robust, sensitive LCMS/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 acidcatalyzed. The rate of perchlorate ion formation was found to be firstorder with respect to hypochlorite and chlorate ions and secondorder 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 21 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. μg/l). However, the results 119

144 revealed that the levels of the chlorate ion are considerably higher than the WHO MCL 11 of 7 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. iii. 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. 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 1113, the decomposition of hypochlorite ion is minimized. 11 WHO Guidelines for Drinking Water Quality 12

145 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 onsite 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 12 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 onsite generation, and thus contribution of the bromate ion to the drinking water will be significantly minimized. 121

146 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 pathways 112 and irreversible oxidation of the reagents/sensing element Puckett, S. D., Heuser, J. A., Keith, J. D., Spendel, W. U. and Pacey, G. E. "Interaction of ozone with gold nanoparticles" Talanta 25, 66, Tomiyasu, H., Fukutomi, H. and Gordon, G. "Kinetics and mechanism of ozone decomposition in basic aqueous solution" Inorg. Chem. 1985, 24, Gordon, Cooper, Rice, and Pacey Disinfectant Residual Measurement Methods. 122

147 Electrochemicallybased 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 watercooled Ozone Research Equipment Corporation Model O3B9 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 UVVis. Molar absorptivity of 33 M 1 cm 1 was used to calculate the concentration of ozone. A 1 cm quartz sample cuvette was used to increase sensitivity of UVVis method. Perkin Elmer, Lambda 95 Reflectance Spectrophotometer, employing an integrating sphere was used to collect absorbance spectra of gold nanoislands. Varian s Cary 5 spectrophotometer was used to collect UVVis 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: 78mtorr vacuum, High Voltage set to generate current above 1 ma and coating time: 1 min5min. 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". 123

148 to produce nanoislands. Lower temperatures require longer annealing time but can be used as well. The synthezid nanoislands was characterized by an Agilent Technologies 55 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 A11 shows a typical SEM image of gold nanoislands on polished aluminum substrate. Figure A12 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 A11. SEM images of typical gold nanoislands produced by sputtering process on a polished aluminum substrate 124

149 (a) (b) Figure A12. AFM images of 25 nm (a); and 14 nm (b) gold nanoislands on quartz substrate As can be seen from Figure A13, gold nanoislands respond to the presence of ozone by producing a redshift in the surface plasmon resonance. The surface plasmon resonance peak for the gold nanoislands shifts from 52 nm to 54 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 125

150 surface plasmon resonance peak for the gold nanoislands completely return to the original position at 52 nm. The top line in Figure A13 shows the spectrum for a thin film of gold that produces a broad absorbance after 6nm 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 reported 117. The oxygen layer could be removed by heating the surface to 6 K and that this desorption process exhibited firstorder kinetics Sputtered gold thin film Gold nanoislands Gold nanoislands exposed to Ozone Gold nanoislands reversed Absorbance Wavelength, nm Figure A13. Overlaid UVVis spectra of : dashed linegold thin film on quarts, solid linegold nanoislands, dotdashed line gold nanoislands exposed to ozone gas, dotted linegold 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, 41, Krozer, A. and Rodahl, M. "Xray photoemission spectroscopy study of UV/ozone oxidation of Au under ultrahigh vacuum conditions" J. Vac. Sc. Technol. A. 1997, 15, Saliba, Parker and Koel "Adsorption of oxygen on Au(111) by exposure to ozone". 126

151 Fortunately, gold nanoislands do not need high temperatures to shift the surface plasmon resonance back to 52 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 surfaceplasmon 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 5. A detection limit of.1 mg/l for ozone with a precision of at least.1% and an accuracy of at least.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 ultracomplex 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 A14 shows the spectra exhibiting the amount of redshift, in nm, with increasing ozone concentration. Figure A15 is a 119 Gordon, Cooper, Rice, and Pacey Disinfectant Residual Measurement Methods. 127

152 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 2 ppb (µg/l). Starting with the 2 ppb data point, a curve with an equation of y=6.8ln(x)15.19 produced a correlation coefficient of Absorbance Wavelength (nm) Figure A14. Overlaid UVVis spectra of 25 nm gold nanoislands with surface Plasmon absorbance max at 52 nm exposed to concentrations of ozone, increased in increments form 2.9 μg/l to μg/l. Ozone causes a redshift in the surfaceplasmon absorbance max 128

153 Shift, nm Ozone Concentration, ppb (ug/l) Figure A15. Shift of the 25 nm gold nanoislands surfaceplasmon max (52 nm) as a function of ozone concentration, logarithmic fit gives an equation of y=6.8ln(x) produced a correlation coefficient of.966 So far the majority of this work is performed with 13 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 A16 shows a plot of surfaceplasmon 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 52 nm. Absorbance max of the surface plasmon suggests that the nanoislands shown in Figure A1 6, are of a larger size. Figure A16 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 129

154 of the particles. It would also make sense to extend the range of nanoislands to sizes between 1 and 1 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 Shift, nm Ozone Concentration, ppb (ug/l) Figure A16. 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 UVVis 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 115 minutes. At 25 ºC 2 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 13

155 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. 131

156 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, 12 iontrack etching, 121 electrochemical etching of Si 122, and anodization of Al 123. 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 surfactanttemplated materials used a liquidcrystal template mechanism in which aluminosilicate gels were calcined in the presence of surfactant 12 Choi, Y.K., King, T.J. and Hu, C. "Nanoscale CMOS spacer FinFET for the terabit era" Electron. Devic. Lett. 22, 23, 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. 24, 14, Sato, H., Homma, T. "Fabrication of highaspectratio arrayed structures using Si electrochemical etching" Sci. Technol. Adv. Mat. 26, 7, Xu, C. X., Zhang, X. S. and Sun, X. W. "Preparation of Porous Alumina by Anodization" J. Metastab. Nanocryst. 25, 23, Sanchez, C., Julian, B., Belleville, P. and Popall, M. J. Mater. Chem. 25, 15, Scott, B. J., Wirnsberger, G. and Stucky, G. D. "Mesoporous and Mesostructured Materials for Optical Applications" Chem. Mater. 21, 13, Hartmann, M. "Ordered Mesoporous Materials for Bioadsorption and Biocatalysis" Chem. Mater. 25, 17, Hoffmann, F., Cornelius, M., Morell, J. and Fröba, M. "SilicaBased Mesoporous OrganicInorganic Hybrid Materials" Angew. Chem., Int. Ed. 26, 45,

157 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 airwater 129 or wateroil systems by using surfactant micelles as a template to grow mesoporous silica at these interfaces. 13 However, it was found that synthesis of mesoporous silica films using surfactant template had long deposition times, the film texture was grainy 131, and many times the orientation of the channels was parallel to the substrate. 132 The next advance was evaporationinduced selfassembly (EISA). 133,134,135,136 This process was based on silica/surfactant selfassembly into thinfilm mesophases during the evaporation of the solvent and of the alcohol released by processing after inkjet printing, spincoating, or dipcoating 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 dipcoating. An overview of synthesis methods that produce mesoporous metaloxide and other oxide films with pore orientation perpendicular to the substrate surface with the use of surfactant templating and dimensional confinement has been presented 137. In particular, a porous alumina membrane can be used to provide a dimensional confinement for the selfassembly 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 liquidcrystal template mechanism" Nature 1992, 359, Yang, H., Coombs, N., Sokolov, I. and Ozin, G. A. "Freestanding and oriented mesoporous silica films grown at the airwater interface" Nature 1996, 381, Schacht, S., Huo, Q., VoigtMartin, I. G., Stucky, G. D. and Schüth, F. "OilWater Interface Templating of Mesoporous Macroscale Structures" Science 1996, 273, 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. 25, 15, Yang, H., Kuperman, A., Coombs, N., MamicheAfara, S. and Ozin, G. A. "Synthesis of oriented films of mesoporous silica on mica" Nature 1996, 379, 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 solgel dipcoating" Nature 1997, 389, Brinker, C. J., Lu, Y., Sellinger, A. and Fan, H. "EvaporationInduced SelfAssembly: Nanostructures Made Easy" Adv. Mat. 1999, 11, Grosso, D., Cagnol, F., SolerIllia, G., thinsp, J, de, A, Crepaldi, E., L, Amenitsch, H., BrunetBruneau, A., Bourgeois, A. and Sanchez, C. "Fundamentals of Mesostructuring Through EvaporationInduced Self Assembly" Adv. Funct. Mater. 24, 14, Minoofar, P. N., Dunn, B. S. and Zink, J. I. "Multiply Doped Nanostructured Silicate SolGel Thin Films: Spatial Segregation of Dopants, Energy Transfer, and Distance Measurements" J. Am. Chem. Soc. 25, 127, Brinker, C. J. and Dunphy, D. R. "Morphological control of surfactanttemplated metal oxide films" Curr. Opin. Colloid In. 26, 11,

158 perpendicular to the membrane s transverse channels. 138,139 However, due to lack of a nonnanostructured 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 dipcoating silica precursors and surfactants onto modified glass slides 14. 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 photocrosslinkable 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 4nm) 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 SBA15 Nanorod Arrays Inside a Porous Alumina Membrane" J. Am. Chem. Soc. 24, 126, Yamaguchi, A., Uejo, F., Yoda, T., Uchida, T., Tanamura, Y., Yamashita, T. and Teramae, N. "Selfassembly of a silicasurfactant nanocomposite in a porous alumina membrane" Nat. Mater. 24, 3, Koganti, V. R. and Rankin, S. E. "Synthesis of SurfactantTemplated Silica Films with Orthogonally Aligned Hexagonal Mesophase" J. Phys. Chem. B 25, 19, 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. 25, 5, Fukumoto, H., Nagano, S., Kawatsuki, N. and Seki, T. "PhotoAlignment Behavior of Mesoporous Silica Thin Films Synthesized on a PhotoCrossLinkable Polymer Film" Chem. Mater. 26, 18, Koganti and Rankin "Synthesis of SurfactantTemplated 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 "PhotoAlignment Behavior of Mesoporous Silica Thin Films Synthesized on a PhotoCrossLinkable Polymer Film". 134

159 Another approach to obtaining mesoporous silica films with normal pore orientation is to combine electrochemically induced selfassembly of surfactants at the solidliquid interface 146 and an electrochemically assisted deposition method to produce solgels. 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 thinfilm growth is guided by both the surfactanttemplated surface and solgel controlled by the potential of the substrate electrode. It is also different from cathodic deposition of ordered metal or metaloxide films around polymeric templates 15 and electrochemical precipitation of metal oxides using selfassembly of surfactant molecules in that the electrogenerated species in this process is a catalyst rather than a component of the final material 151. 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, solgel 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 selfassembly of surfactants at the solidliquid 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 PotentialControlled Surfactant Assembly" Adv. Mat. 23, 15, Shacham, R., Avnir, D. and Mandler, D. "Electrodeposition of Methylated SolGel Films on Conducting Surfaces" Adv. Mat. 1999, 11, 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 26, 22, Walcarius, A., Sibottier, E., Etienne, M. and Ghanbaja, J. "Electrochemically assisted selfassembly of mesoporous silica thin films" Nat. Mater. 27, 6, 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. 2, 2, Choi, K.S., Lichtenegger, H. C., Stucky, G. D. and McFarl, E. W. "Electrochemical Synthesis of Nanostructured ZnO Films Utilizing SelfAssembly of Surfactant Molecules at Solid Liquid Interfaces" J. Am. Chem. Soc. 22, 124,

160 method has been previously reported 152. 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 solgel 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 solgel is also perpendicular to the substrate surface. The CTAB is leached from the film, resulting in a solgel coated surface in which the perpendicular mesoporous structure is maintained. In the reported study 153, 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. A22. Experimental The organically modified precursor was prepared by reacting 7hydroxy4methyl coumarin (coumarin 4) with 3(triethoxysilyl) propylisocyanate (TriMeOSiC) to produce 4methylcoumarin7yl 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; 2 ml ethanol; 2 ml aqueous solution of.1 M NaNO3; 13 M HCl, and an aliquot of 13 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.32 (higher or lower ratios give less ordered mesostructures or even nonordered 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 1 s, depending on the desired film thickness). 152 Walcarius, Sibottier, Etienne and Ghanbaja "Electrochemically assisted selfassembly of mesoporous silica thin films". 153 Walcarius "Electrochemically assisted selfassembly of mesoporous silica thin films". 154 Walcarius "Electrochemically assisted selfassembly of mesoporous silica thin films". 136

161 Under these conditions the reduction of water occurs; the OH released catalyzes the solgel 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 evaporationinduced 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 11 C. The reported drying at 13 C 155 was not used to avoid possible thermal decomposition of 4methylcoumarin7yl 3 trimethoxysilyl)propylcarbamate. The dried samples were extracted overnight with 1 ml of 5:5 ethanol.1m HCl (aq.) to extract CTAB as well as residual sol. All fluorescence measurements were obtained with a Horiba Jobin Yvon Fluorolog322 spectrofluorometer. AFM images were obtained with a Agilent 55 AFM Microscope. A23. Results and Discussion An initial screening of the fluorescence from films on ITO is shown in Figure A2 1. Under UV radiation (Fig. A21b), fluorescence is observed. In addition, the EPONcoated ITO appeared white due to light scattering. Note that the sample used in Figure A21 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 solgel structures. (a) (b) Figure A21. (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 selfassembly of mesoporous silica thin films". 137

162 AFM images of the EPON on ITO are shown in Figure A22. Here the electrochemically assisted deposition was 1 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 nonuniform catalysis, so the electrolysis of water, and therefore the ph, will vary across the electrode surface, as shown by Figure A22. 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 A23. The EPON pore structure previously reported with a TEOS precursor 156 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 A22. EPONCoated ITO Electrode, plating time 1 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 selfassembly of mesoporous silica thin films". 138

163 (a) (b) Figure A23. EPONcoated ITO electrode, plating time: (a) 3 s 2.1V. The deposited film indicates normal to the electrode surface orientation of the deposited EPON film. (b) EPONcoated ITO electrode, magnification of (a) reveals normal orientation of the mesopores The UVVis and fluorescence spectra of 7hydroxy4methyl coumarin and 4 methylcoumarin7yl 3 (trimethoxysilyl)propylcarbamate solutions of the same concentration (1 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 siloxyderivatized coumarin 4 that was incorporated into solgel 157,158, 159 monoliths. The incorporation of coumarin 4 in the solgel film is stable after the initial washing to remove the unbound fluorophore. For example, Figure A24 compares the fluorescence from a film at various stages of fabrication. After washing the coumarin4 emission drops by a factor of 2.2, indicating the presence of nonbonded 4 methylcoumarin7yl 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 SolGel Hosts. 1. Structure and Environmental Factors on Fluorescent Properties" Chem. Mater. 1998, 1, Suratwala, T., Gardlund, Z., Davidson, K., Uhlmann, D. R., Watson, J., Bonilla, S. and Peyghambarian, N. "Silylated Coumarin Dyes in SolGel Hosts. 2. Photostability and SolGel Processing" Chem. Mater. 1998, 1, Suratwala, T., Gardlund, Z., Boulton, J. M. and Uhlmann, D. R. Series Incorporation of triethoxysilyl functionalized Coumarin 4 in solgel hosts, 13 October, 1994 Conference. 139