MIAMI UNIVERSITY
The Graduate School
Certificate for Approving the Dissertation
We hereby approve the Dissertation
of
Zsolt Körtvélyesi
Candidate for the Degree:
Doctor of Philosophy
Director
Gilbert Gordon
Chair
Gilbert E. Pacey
Reader
James A. Cox
Reader
Michael W. Crowder
Graduate School Representative
John M. Hughes
ABSTRACT
ANALYTICAL METHODS FOR THE MEASUREMENT OF CHLORINE DIOXIDE AND
RELATED OXYCHLORINE SPECIES IN AQUEOUS SOLUTION
by Zsolt Körtvélyesi
The main goal of this research was to seek a better understanding of the analytical measurements
of the oxychlorine species ClO2, Cl2O4–, and Cl2O3/Cl2O3–.
The US EPA has developed a new colorimetric method for the measurement of ClO2 and chlorite
ion (Method 327.0). This method is based on the decolorization of the dye Lissamine Green B
(LGB) by ClO2. Chlorite ion is converted to ClO2 by Horseradish Peroxidase enzyme and
measured with LGB. In the current work, the performance of this method (method detection
limit, accuracy, and precision) was evaluated. The interference from dissolved chlorine,
chloramine, iron(II), manganese(II), permanganate, and chlorate ions was studied. The underlying
chemistry of these reactions is described and used to differentiate between interference and
demand. A new method is suggested for the preparation of ClO2 standards by illuminating a
mixture of chlorite ion and a photoacid. By using this method, ClO2 standards could be prepared
reproducibly. Possible future developments for the method are also discussed.
Chlorite ion interferes with the spectrophotometric measurement of ClO2 due to the formation of
the Cl2O4– complex. This complex has higher molar absorptivity than ClO2 at longer wavelengths
where the absorbance of concentrated ClO2 solutions is measured. The formation constant of the
complex is 5.0 M–1 as determined in this work. Based on this value, the molar absorptivity of the
complex was calculated as a function of wavelength. These values were used to give recommendations to adjust the currently used spectrophotometric measurements.
A new mixed disinfectant solution was developed and tested. This disinfectant is created from
dissolved chlorine and ClO2. It is a potent disinfectant due to the formation of reactive
intermediates resulting from the ClO2–chlorine reaction. A combination of chemical kinetic and
microbiological results was used to estimate the efficacy of the new solutions. It was shown that
in this way, fewer microbiological tests are required than using only microbiological results. This
leads to shorter development time and lower costs.
ANALYTICAL METHODS FOR THE MEASUREMENT OF CHLORINE DIOXIDE AND
RELATED OXYCHLORINE SPECIES IN AQUEOUS SOLUTION
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
Zsolt Körtvélyesi
Miami University
Oxford, Ohio
2004
Dissertation Advisor: Dr. Gilbert Gordon
©
Zsolt Körtvélyesi
2004
T ab le of co ntents
List of tables
viii
List of figures
x
List of Acronyms
xv
Acknowledgments
xvii
1. Introduction and Research Objectives
1
1.1. The chemistry of chlorine in aqueous solutions
2
1.2. Chlorine dioxide
3
1.2.1. Physical, chemical properties of ClO2
3
1.2.2. Properties of sodium chlorite
5
1.2.3. Generation of ClO2
6
1.2.4. Applications of ClO2 in water treatment
8
1.2.5. Safety precautions for research laboratories
1.3. Research objectives
10
10
2. Proposed EPA Method 327.0: Determination of ClO2 and Chlorite Ion in Drinking Water Using
Lissamine Green B and Horseradish Peroxidase (HRP) with Detection by Visible
Spectrophotometry
12
2.1. Regulations of ClO2 in potable water
13
2.2. Current ClO2 analytical methods
13
2.2.1. Ideal Method
14
2.2.2. Iodometric method
15
2.2.3. Spectrophotometric method
16
2.2.4. Colorimetric methods
19
2.2.5. N, N’-diethyl-p-phenylenediamine (DPD) method
20
iii
2.2.6. Lissamine Green B (LGB) method
23
2.2.7. Other colorimetric methods
24
2.2.8. Electrochemical methods
25
2.3. Experimental
26
2.3.1. Reagent water
26
2.3.2. Generation of ClO2
26
2.3.3. Carbonate free sodium hydroxide solutions
27
2.3.4. Preparation of dissolved chlorine solutions
27
2.3.5. Preparation of monochloramine
27
2.3.6. Titration of chlorine and monochloramine solutions
28
2.3.7. Titration of ClO2 and chlorite ion solutions
28
2.3.8. Preparation and titration of potassium permanganate solutions
29
2.3.9. Other reagents
29
2.3.10. Shrinking bottle
29
2.3.11. Other equipment
30
2.4. The proposed LGB method
31
2.5. Results of the second laboratory experiments
33
2.5.1. Method detection limit
37
2.5.2. Recoveries of the samples
38
2.6. Interference studies
41
2.6.1. Interference in analytical measurements
41
2.6.2. Demand vs. interference
43
2.6.3. The interferences studied
45
2.6.4. Chlorate ion interference
46
2.6.5. Iron(II) interference
48
2.6.6. Manganese(II) interference
58
2.6.7. Manganese(VII) interference
61
2.6.8. Manganese(II)–Manganese(VII) interference
65
2.6.9. Free Available Chlorine (FAC) interference
71
iv
2.6.10. Monochloramine (NH2Cl) interference
73
2.6.11. Conclusions on the interference results
74
2.7. Conclusions
77
2.8. Future directions
78
2.8.1. Chlorine dioxide standards
78
2.8.2. Using gas diffusion flow injection analysis with proposed EPA Method 327.0
80
3. The Cl2O4– Complex
82
3.1. Theoretical
83
3.1.1. The history of the Cl2O4– complex
3.2. Numerical methods
83
85
3.2.1. Matrix Rank Analysis
85
3.2.2. Determination of formation constants
89
3.2.3. PSEQUAD
91
3.2.4. Excel workbook for the determination of formation constants
93
3.3. Experimental
94
3.3.1. Purification of sodium chlorite
94
3.3.2. Preparation of sodium perchlorate solution
95
3.3.3. Determination of the molar absorptivity of ClO2 and chlorite ion
96
3.4. Problems with the spectrophotometric measurement of the Cl2O4– complex
98
3.5. Long-period grating (LPG) sensor results
98
3.5.1. The calibration of the LPG sensor for the determination of chlorite ion concentration
102
3.5.2. The calibration of the LPG sensor for the determination of ClO2 concentration
106
3.5.3. The response of the LPG sensor in mixtures of ClO2 and chlorite ion
107
3.6. Initial spectrophotometric results
108
3.7. Main spectrophotometric study
112
3.7.1. The effect of temperature on the equilibrium
v
118
3.8. The structure of the complex
119
3.9. Methods to eliminate the interference of the complex
123
3.9.1. Comparison of Cl2O4– complex with other similar species
3.10. Conclusions
126
127
4. Measurement of Reactive Species and Intermediates in Mixed Disinfectant Solutio ns: The
Dissolved Chlorine (Free Available Chlorine, FAC)–ClO2 System
129
4.1. Theoretical
130
4.1.1. The FAC-ClO2 reaction
130
4.1.2. The C×T principle
132
4.1.3. Microbiological tests
133
4.2. Experimental
135
4.2.1. Preparation of the disinfectant (ClO2, FAC) solutions
135
4.2.2. Analytical methods for the measurement of the various species
135
4.2.3. Iodometric titration
136
4.2.4. Spectrophotometric measurement
138
4.3. Preparation of the Mixed Disinfectant Solutions
140
4.3.1. Generation of ClO2 by mixing FAC and chlorite ion
143
4.3.2. Generation of ClO2 by mixing sodium chlorite with strong acid
148
4.3.3. The reaction of chlorite ion with sodium bisulfate
149
4.3.4. The reaction of chlorite ion with hydrochloric acid
149
4.3.5. Preparation of FAC solutions
151
4.3.6. Preparation of phosphate buffers
151
4.3.7. Preparation of mixed disinfectant solutions
151
4.4. Initial experiments
152
4.5. Kinetic study
154
4.5.1. Temperature effect
161
4.6. Microbiological results
162
4.6.1. Second set of microbiological studies
vi
166
4.7. The effect of the penicylinders on the FAC–ClO2 reaction
4.7.1. Results of iodometric measurements
4.8. Conclusions
167
169
172
5. References
174
Appendix A
183
A.1 Program for collecting data from a Radiometer autotitrator
183
A.2 Program for converting raw data from an Applied Photophysic SF to Excel format
188
vii
List of tables
Table 1. The physical properties of ClO2.
4
Table 2. The electrode potentials of ClO2 at various pH values.
5
Table 3. Demonstration of the effect of the resolution on the molar absorptivity
17
Table 4. Summary of various colorimetric methods used for the measurement of ClO2.
21
Table 5. Parameters of the calibration curves on various days.
34
Table 6. Comparison of the calibration curves for the LGB method by using multiple linear model
regression.
36
Table 7. Detection limits for chlorite ion and ClO2 in reagent water.
37
Table 8. Determined ClO2 concentrations in the presence of chlorate ion, in mg/L units.
47
Table 9. Determined ClO2 concentrations in the presence of iron(II) ion, in mg/L units.
54
Table 10. Determined ClO2 concentrations in the presence of Mn(II), in mg/L units
59
Table 11. Determined ClO2 concentrations in the presence of permanganate ion, in mg/L units.
62
Table 12. Determined ClO2 concentrations in the presence of permanganate ion and
manganese(II), in mg/L units.
67
Table 13. Determined ClO2 concentrations in the presence of FAC, in mg/L units.
72
Table 14. Determined ClO2 concentrations in the presence of monochloramine, in mg/L units.
75
Table 15. Summary of the measured ClO2 concentrations in the presence of various interferents,
in mg/L units. No ClO2 was added to the solutions.
77
Table 16. Repeatability of the wavelength of the valley in water for sensor #15.
101
Table 17. Comparison of the position of the valley for three sensors.
101
Table 18. Results of the calibration of LPG sensor for chlorite ion.
102
Table 19. Comparison of the calibration curves for the same sensor on different days.
103
Table 20. Parameters of calibration curves for chlorite ion determination for various sensors
104
viii
Table 21. The effect of the Cl2O4– complex on the spectrophotometric measurement of ClO2.
c(ClO2)calculated = A450 nm/e450 nm(ClO2), % error = [c(ClO2)calculated–c(ClO2)]/c(ClO2)×100
110
Table 22. Results of an MRA run. Number of spectra = 78, wavelength range: 395–600 nm
111
Table 23. Molar absorptivity of ClO2, NaClO2, and Cl2O4–
117
Table 24. The change of the equilibrium constant with temperature.
119
Table 25. The use of Equation 54 to correct for the presence of the Cl2O4– complex.
c(ClO2)calculated = A449.7 nm/e449.7(ClO2), % error = [c(ClO2)calculated - c(ClO2)]/c(ClO2)×100
125
Table 26. Comparison of the molar absorptivities of the various chlorine containing species at the
wavelengths of the maximum absorptivities (M–1cm–1). Bold numbers indicate the
maximum molar absorptivity for the given species.
139
Table 27. Comparison of the order of FAC at various constant ClO2 concentrations and pH
values.
158
Table 28. Comparison of the order of ClO2 at various constant FAC concentrations and pH
values.
159
Table 29. Comparison of the rate constants for the different reaction pathways and temperatures.
160
Table 30. The composition of mixed disinfectant solutions for the second microbiological studies.
165
Table 31. Comparison of the concentration change of ClO2 in the absence and in the presence of
a penicylinder. Initial concentrations: [FAC] = 300 mg/L, [ClO2] = 300 mg/L, pH 7.0
168
ix
List of figures
Figure 1. The distribution of various FAC species as a function of pH. — Cl2, — HOCl, — OCl–
3
Figure 2. The change in the measured absorbance of a ClO2 solution as a function of the
resolution. [ClO2 ] = 0.0395 M, path length = 0.0098 cm. — 0.1 nm resolution, — 0.5
nm resolution, — 1.0 nm resolution, — 2.0 nm resolution
17
Figure 3. Drawing of a shrinking bottle. A - precision screw, B - brass frame, C - retaining
springs, D - guide for screw, E - 50 mL syringe
Figure 4. The chemical structure of LGB
30
32
Figure 5. The percent recoveries of pure chlorite ion standards, which were determined during
the second laboratory testing. — 0.25 mg/L chlorite ion,  1.0 mg/L chlorite ion, •
2.0 mg/L chlorite ion.
39
Figure 6. The percent recoveries of pure ClO2 standards which were determined during the
second laboratory testing. — 0.25 mg/L ClO2,  0.8 mg/L ClO2, • 2.0 mg/L ClO2. 39
Figure 7. The change of the determined ClO2 concentration with chlorate ion concentration.  no
ClO2 added, Ž 0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added. The dashed lines show
0, 0.25, and 2.0 mg/L determined ClO2 concentrations.
46
Figure 8. Measured spectra of the LGB solution after the addition of ClO2 solutions. The inset
shows the spectral region of 275 nm to 325 nm. — No ClO2, — 0.5 mg/L ClO2, — 1.0
mg/L ClO2, — 2.0 mg/L ClO2
49
Figure 9. Measured spectra of the LGB solution after the addition of iron(II) solution in the
absence of ClO2. The inset shows the spectral region 275 nm to 325 nm. — No Fe(II),
— 1.00 mg/L Fe(II), — 2.00 mg/L Fe(II), — 5.00 mg/L Fe(II), — 9.99 mg/L Fe(II)
50
Figure 10. The absorbance change at 303 nm as the function of Fe(II) concentration. The line is
the least squares fitted line. The equation of this line: Abs303 nm = 0.0310×[Fe2+] +
0.255, R2 = 0.999
51
x
Figure 11. Measured spectra of the LGB solution after the addition of Fe(II) solution in the
absence of glycine/citric acid buffer and ClO2. The inset shows the spectral region 275
nm to 325 nm. — No Fe(II), — 1.00 mg/L Fe(II), — 2.00 mg/L Fe(II), — 5.00 mg/L
Fe(II), — 9.99 mg/L Fe(II)
52
Figure 12. The change of the determined ClO2 concentration with iron(II) concentration.  no
ClO2 added, Ž 0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added. The dashed line
show 0.25 mg/L determined ClO2 concentration.
53
Figure 13. The change of the determined ClO2 concentration with Mn(II) concentration.  No
ClO2 added, Ž 0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added
59
Figure 14. The change of the determined ClO2 concentration with permanganate ion
concentration. The lines show the least square fit of the data.  No ClO2 added, Ž
0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added. The equation of the lines: —
[ClO2]det. = 0.245×[MnO4–] + 0.464, R2 = 0.970, — [ClO2]det. = 0.270×[MnO4–] +
0.402, R2 = 0.972
62
Figure 15. The change of the determined ClO2 concentration with the permanganate ion to ClO2
ratio.
64
Figure 16. The change of the determined ClO2 concentration with permanganate ion
concentration in the presence of Mn(II). No ClO2 was added.  0.1 mg/L Mn(II), Ž
1.0 mg/L Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
68
Figure 17. The change of the determined ClO2 concentration with the permanganate ion
concentration in the presence of Mn(II). 0.25 mg/L ClO2 was added.  0.1 mg/L
Mn(II), Ž 1.0 mg/L Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
68
Figure 18. The change of the determined ClO2 concentration with permanganate ion
concentration in the presence of Mn(II). 2.0 mg/L ClO2 added.  0.1 mg/L Mn(II), Ž
1.0 mg/L Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
69
Figure 19. The change of the determined ClO2 concentration with the manganese(II) to
permanganate ion molar ratio.  1.0 mg/L MnO4–, Ž 2.0 mg/L MnO4–, • 5.0 mg/L
MnO4–, – 10.0 mg/L MnO4–
70
xi
Figure 20. The change of the determined ClO2 concentration with FAC concentration.  no ClO2
added, Ž 0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added
71
Figure 21. The change of the determined ClO2 concentration with monochloramine
concentration.  no ClO2 added, Ž 0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added
75
Figure 22. The molar absorptivities of ClO2 and chlorite ion. The inset shows the 320 nm to 450
nm region. — Sodium chlorite, — ClO2
97
Figure 23. Signal of an LPG sensor in air.
100
Figure 24. Signal of an LPG sensor in water
100
Figure 25. Calibration curve of an LPG sensor for chlorite ion. The equation of the line is λvalley =
2.95×[NaClO2] + 824.61, R2=0.965
103
Figure 26. Calibration curves of an LPG sensor for chlorite ion using the same fiber on different
days. For the parameters of the calibration curves see Table 20.  Day 1, Ž day 2, •
day 3
104
Figure 27. Calibration curves of various LPG sensors using the same chlorite ion solutions and
different fibers. For the parameters of the calibration curves see Table 21.  #3, Ž
#15, • #16
105
Figure 28 Calibration curve of an LPG sensor for ClO2.
106
Figure 29. Response of an LPG sensor at different Cl2O4– complex concentrations. The
concentration of the complex was determined by using Keq = 1.6 M–1.
Figure 30. a)Photograph and b) schematic drawing of tandem cell
107
108
Figure 31. Absorbance change of a ClO2 and chlorite ion mixture before (—) and after mixing
(þ). c(ClO2) = 168.4 mg/L (2.5×10–3 M), c(NaClO2) = 112.0 g/l (1.66 M), A450 nm =
0.130 before mixing, A450 nm = 0.184 after mixing, path length = 2×0.437 cm
109
Figure 32 Residual spectra after assuming — 1, — 2, — 3, and — 4 absorbing species.
112
Figure 33. Results of fitting of stopped-flow data by using PSEQUAD. Temperature = 25 °C
114
Figure 34. Results of fitting of stopped-flow data by using Excel worksheet. Temperature = 25 °C
115
xii
Figure 35. Comparison of the molar absorptivity of the various species in the chlorite ion–ClO2
system.
116
Figure 36. The change of the formation constant of the Cl2O4– complex with temperature. The
equation of the least squares fitted line is log Keq = 2293.8×T–1 – 7.06, R2 = 0.958 119
Figure 37. EPR spectra of ClO2 and ClO2/chlorite ion mixture at room temperature. [ClO2] =
2.46×10–3 M in ClO2 solution (—), [ClO2] = 2.46×10–3 M, [ClO2–] = 5.07×10–2 M in
Cl2O4– solution (—). The EPR spectra were collected with a center field of 3370 G,
sweep width of 200 G, a microwave frequency of 9.424 GHz, modulation frequencey
of 100 kHz, modulation amplitude of 10 G, and a power of 0.635 mW.
Figure 38. The possible structures of the Cl2O4– complex.
120
121
Figure 39. Comparison of the molar absorptivities of the various chlorine containing species in
the mixed disinfectant system. — Hypochlorous acid, — chlorite ion, — hypochlorite
ion, — ClO2
139
Figure 40. The formation of ClO2 with excess FAC as a function of time at various pH values.
[FAC] = 1.07×10–2 M, [ClO2–] =5.97×10–3 M, pH =  7.0, Ž 7.5, • 8.0
145
Figure 41. The formation of ClO2 with excess chlorite ion as a function of time at various pH
values. [FAC] = 1.07×10–2 M, [ ClO2–] = 1.79×10–2 M, pH =  7.0, Ž 7.5, • 8.0 145
Figure 42. The formation of ClO2 with excess chlorite ion as a function of time. The chlorite ion
concentrations are adjusted to reach similar ClO2 concentrations. At pH 7.0 (),
[FAC] = 1.07×10–2 M, [ClO2–] = 1.79×10–2 M; at pH 7.5 (Ž), [FAC] = 1.06×10–2 M,
[ClO2–] = 2.99×10–2 M
146
Figure 43. The dependance of the concentration of the generated ClO2 on the chlorite ion
concentration at constant FAC concentration. pH = 7.0, FAC = 1.07×10–2 M, 
Maximum ClO2, Ž after 5 minutes, • after 10 minutes. The equation of the least
squares fit for the maximum ClO2 concentration: c(ClO2, M)max = 0.14×c(ClO2–, M) –
2.4×10–4
147
Figure 44. The dependence of the ClO2 concentration generated on the initial chlorite ion
concentration. c(HCl) = 0.5 M The equation of the linear fit: c(ClO2, mol/L) =
0.36×c(NaClO2–, mol/L) + 2.7×10–4
150
xiii
Figure 45. Concentration changes of the main species in the FAC–ClO2 system. [ClO2] =
1.48×10–3 M (100 mg/L), [FAC] = 2.12×10–3 M (150 mg/L), pH = 7.5, temperature =
22 °C.  Chlorite ion, Ž ClO2, • FAC
155
Figure 46. Determination of the reaction order of FAC at pH 6.5 by using the method of initial
rates.
157
Figure 47. Determination of the reaction order of FAC at pH 7.5 by using the method of initial
rates.
157
Figure 48. Determination of the reaction order of ClO2 at pH 6.5 by using the method of initial
rates.
158
Figure 49. Determination of the reaction order of ClO2 at pH 7.5 by using the method of initial
rates.
159
Figure 50. Comparison of the measured and fitted ClO2 and FAC concentrations.  Measured
FAC concentration, Ž Measured ClO2 concentration, the solid lines represent the
fitted concentrations
161
Figure 51. Comparison of the ClO2 concentration change in the absence (—) and in the presence
(—) of a penicylinder. Initial concentrations: [FAC] = 300 mg/L, [ClO2] = 300 mg/L,
pH 7.0. See Table 31 for ClO2 concentrations at a given time.
168
Figure 52. The concentration change of FAC in a FAC–ClO2 mixture. [FAC]0 = 300 mg/L,
[ClO2]0 = 300 mg/L, pH = 7.0.  predicted values, Ž in the absence of a penicylinder,
• in the presence of a penicylinder
170
Figure 53. The concentration change of ClO2 in a FAC–ClO2 mixture. [FAC]0 = 300 mg/L,
[ClO2]0 = 300 mg/L, pH = 7.0.  predicted values, Ž in the absence of a penicylinder,
• in the presence of a penicylinder
170
xiv
List of Acronyms
ACVK
Acid chrome violet K
AWWA
American Waterworks Association
CCCS
Continuing Calibration Check Standard
CPR
Chlorophenol red
DBP
Disinfection Byproduct
D/DBPR
Disinfectants and Disinfection Byproducts Rule
DPD
N,N’-diethyl-p-phenylenediamine
EDTA
Ethylenediamine tetraacetic acid
EPA
Environmental Protection Agency
FAC
Free available chlorine: elemental chlorine, hypochlorous acid, hypochlorite ion
HAA5
Haloacetic acids
HRP
Horseradish Peroxidase
LGB
Lissamine Green B
LPG
Long-period grating (sensor)
MCL
Maximum Contaminant Level
MRA
Matrix Rank Analysis
MRDL
Maximum Residual Disinfectant Level
PCA
Principal Component Analysis
PDA
Photodiode array (detector)
xv
PSEQUAD
Potentiometric and/or Spectrophotometric Equilibrium Data Using Analytical
Derivatives
SDWA
Safe Drinking Water Act
SF
Stopped-flow (spectrophotometer)
TDW
Triple-distilled water
THM
Trihalomethane
VBA
Visual Basic for Applications
xvi
Acknowledgments
I would like to thank my research director, Professor Gordon, for all his help during my graduate
career. He taught me not to settle for the almost perfect, but go for the best. He has been a great
mentor, I have learned many things from him and not just about chemistry.
I would like to thank Dr. Pacey, Sean Puckett, and Jason Keith for their help in using the long
period grating sensors. Dr. Crowder and Jim Garety assisted me in using the stopped-flow. I would
also like to thank Kyle Ellis for his contributions to the LGB work. I thank Dr. Gábor Lente for his
help in using the programs MRA and PSEQUAD.
A great thank you goes to Ágnes Balogh, who supported me during the hard times and gave me
strength not to give up. In addition, she aided me in correctly using the various statistical methods
and patiently answered my questions.
I also thank Jennifer Anderson, Meghan Holdorf, Nathan Wenzel, Matt Breece, Jason Keith,
Sean Puckett, Sawmya Chandrasekar, and Craig Gibeau for making my time at Miami a pleasant
experience.
xvii
1. Introduction and Research Objectives
It was recognized by the end of the nineteenth century that many illnesses are caused by
waterborne microorganisms1, 2. In North America aqueous chlorine was the first disinfectant that
gained wide use in water treatment1, 2. During the twentieth century other disinfectants were
discovered and tested for the disinfection of potable water. In 1974, chloroform (a possible
carcinogen) was discovered in chlorinated potable water3. This chemical is formed in the reaction of
dissolved chlorine with natural organic matter in water. Following this discovery, intensive research
was conducted to determine the reaction products of chlorination of naturally occurring organic
materials. This research identified further chlorinated and probable carcinogenic organic compounds
in chlorinated water. These halogenated organic compounds include2, 4 trihalomethanes (THMs) and
haloacetic acids (HAA5). Research was conducted to establish the safe levels of these compounds,
which do not result in adverse effects. Based on these “safe” levels, several regulations came into
effect, e.g., Safe Drinking Water Act (SDWA)5, Disinfectants and Disinfection Byproducts Rule
(D/DBPR)6, 7.
The discovery of toxic chemicals in chlorinated water and the resulting regulations led to an
increased interest in alternative disinfectants that may be able replace dissolved chlorine in water
treatment applications. The two main alternatives are ClO2 and ozone2, 4. In the last thirty-forty years,
intensive research has been conducted with these disinfectants in order to obtain a better
understanding of their chemical properties and their effects on various microorganisms.
1
1.1. The chemistry of chlorine in aqueous solutions
Chlorine is soluble in water2, 4 and its dissolution in water is a physical and chemical process.
Upon contact with water, chlorine disproportionates2, 8, giving chloride ion and hypochlorous acid
or hypochlorite ion, depending on the pH of the solution. Thus, dissolved chlorine is present in
aqueous solutions as a mixture of three active species, molecular chlorine (Cl2), hypochlorous acid
(HOCl), and hypochlorite ion (OCl–). These three species are generally called free available chlorine9
(FAC). The distribution of these species can be described* by the following equations8:
Cl2 + H2O ¾ Cl– + H+ + HOCl
KH = 4×10–4 (25°C)
(1)
HOCl ¾ H+ + OCl–
pKa = 7.54 (25°C)
(2)
Both equilibria are pH dependent. Thus, the distribution of the three species varies with the pH
of the solutions. By using the equilibrium constants, a distribution diagram (Figure 1) can be created
that shows the relative abundance of each of these species at various pH values. Figure 1shows that
below pH 2, molecular chlorine is the predominant species. Above pH 9, FAC exists almost entirely
in the form of hypochlorite ion. These facts are very important in interpreting the properties and the
reaction mechanisms of dissolved chlorine.
It is well-known that molecular chlorine and hypochlorous acid are more effective biocides than
hypochlorite ion2, 4. Increasing the pH above 7 results in a decrease in the disinfectant efficacy of
dissolved chlorine due to the increasing concentration of hypochlorite ion in this region. Therefore,
*
In this thesis the following convention is used to denote the reversibility of the reactions. If the reaction is reversible
and takes place in one step, the double arrow (¾) is used. If the reaction is reversible but takes place through a complex
reaction mechanism, the equal sign is used. If the reaction is irreversible, a single arrow (Ž) is used.
2
keeping the pH of the disinfected water below this value is important to maximize the disinfection
efficacy of FAC.
Figure 1. The distribution of various FAC species as a
function of pH. — Cl2, — HOCl, — OCl–
The advantage of FAC is that it is the most widely used disinfectant in North American and a
significant amount of practical knowledge has been accumulated about its application. Dissolved
chlorine is the easiest and least expensive form of disinfection. Furthermore, dissolved chlorine
provides residual in the treated water. The main disadvantage of FAC is that it forms halogenated
disinfection byproducts.
1.2. Chlorine dioxide
1.2.1. Physical, chemical properties of ClO 2
Chlorine dioxide (ClO2) is a greenish yellow gas at room temperature with an odor that resembles
chlorine2, 9-11. Chlorine dioxide is a very reactive species. Above –40°C it is unstable in pure form4, 12.
3
It undergoes explosive decomposition if its concentration exceeds 10% by volume in air4, 11.
Concentrated solutions present a problem if the ClO2 partial pressure exceeds 10.1 kPa.
One of the most interesting properties of ClO2 is that it exists as a stable free radical12, 13 and no
dimerization reactions occur under normal circumstances. The O–Cl–O bond angle12 is 117.5°, and
the chlorine–oxygen bond length is 1.47 D. This bond has a double bond character. Some parameters
of ClO2 are summarized in Table 1.
Table 1. The physical properties of ClO2.
Molecular weight (g/mole)
67.45
Melting point (°C)
–59
Boiling point (°C)
11
Dipole moment (Debye)
1.69
Henry constant14 (M/atm)
1.0
Chlorine dioxide readily dissolves in water2, 4, 11, but unlike chlorine, ClO2 does not undergo a
chemical reaction. Chlorine dioxide also dissolves in organic solvents, e.g., carbon tetrachloride.
Despite its high solubility in water, ClO2 is volatile. It is important to keep this fact in mind when ClO2
solutions are used. If the necessary precautions are taken, evaporation loss of ClO2 can be minimized.
Neutral or acidic aqueous solutions of ClO2 are stable for long periods of time if they are stored in
the dark, at cool temperatures with no headspace2, 4, 11.
Chlorine dioxide is a strong but selective oxidizing agent. The electrode potentials of various
reactions, which include ClO2, are given in Table 2. In most of its reactions, it undergoes a one
electron transfer reaction forming chlorite ion. Under appropriate conditions, chlorite ion can react
further forming chloride ion. When ClO2 reacts with organic molecules, the carbon-carbon bonds are
4
generally not cleaved and no addition of chlorine to the organic molecule occurs. Thus, no chlorinated
organic compounds are formed12. Chlorine dioxide reacts with phenolic compounds12, and as a result,
ClO2 is very effective in removing phenolic tastes and odors from treated water. Chlorine dioxide
reacts rapidly with organic sulfides and tertiary amines15, 16. The reactions between ClO2 and primary
and secondary amines15, 16, alcohols, and carbonyl compounds are slow. Chlorine dioxide reacts with
many inorganic species4, 11, including manganese(II), iron(II), and aqueous chlorine17, 18.
Table 2. The electrode potentials of ClO2 at various pH values19.
ClO2(g) + 1e– Ë ClO2–
ClO2(g) + H+ + 1e– Ë
HClO2
ClO2(g) + 4 H+ + 5e– Ë
Cl– + 2 H2O
Chlorine dioxide undergoes various disproportionation and self-decomposition reactions12. It
disproportionates to chlorite and chlorate ions in basic solutions12, 20.
2 ClO2 + 2 OH– = ClO2– + ClO3– + H2O
(3)
The reaction is relatively slow, but around pH 9 it consumes ClO2 rapidly. Chlorine dioxide
undergoes thermal decomposition and photochemical decomposition reactions.
1.2.2. Properties of sodium chlorite
Sodium chlorite is a white crystalline material in its pure form21. It is a highly reactive, strong
oxidant. Chlorine dioxide is generated when sodium chlorite comes into contact with acids or
5
chlorine2,
10, 11
(either gaseous or dissolved chlorine). Sodium chlorite violently reacts with
combustible materials21.
Sodium chlorite is available as technical grade solid, containing ~80% (m/m) sodium chlorite10, 21.
The other major components of this technical grade chemical are sodium chloride, sodium carbonate,
and sodium hydroxide. For this reason, solutions of technical grade sodium chlorite are basic. Sodium
chlorite is also available as dilute aqueous solution with up to 40% (m/m) sodium chlorite content.
Sodium chlorite readily dissolves in water. In the case of pure sodium chlorite, the color of this
solution changes from colorless to pale yellow, depending on the chlorite ion concentration. This
color is due to the absorption peak of chlorite ion in the UV region that tails into the visible region.
The maximum absorbance is at 260 nm, but concentrated solutions can absorb light significantly even
around 380 nm.
Chlorite ion solutions are relatively stable if protected from light. However upon exposure to
light, rapid decomposition takes place. The products of this decomposition include chlorate, chloride
ions, oxygen, and possibly ClO2. Sodium chlorite reacts with acids and chlorine, producing ClO2 as
described in the next section.
1.2.3. Generation of ClO 2
According to the DOT regulations22, ClO2 can not be transported. Thus, it needs to be generated
at the point of use. Extensive reviews have been published on the various generation methods and the
details of the generators. Here, only the general chemical considerations are reviewed. Specific
technical details about the generation methods can be found elsewhere2, 4, 10, 11, 23.
The two chemicals, which are used for ClO2 generation, are sodium chlorite and sodium chlorate.
Traditionally, sodium chlorite was the source of the ClO2 generated in water treatment applications.
6
However, the situation is changing and currently there are sodium chlorate-based ClO2 generators
for water treatment purposes.
Chlorite ion – chlorine system: Chlorite ion can be oxidized to ClO2 by using chlorine. The
reaction is described9 by Equation 4.
2 ClO2– + Cl2(g) Ž 2 ClO2 + 2 Cl–
(4 a)
2 ClO2– + HOCl Ž 2 ClO2 + Cl– + OH–
(4 b)
However, these equations give a simplistic representation of the generation process. Considering
the mechanism of these reactions is important for a better understanding of the details of the
generation process. The intermediate species, Cl2O2, forms in these reactions. This intermediate may
react further to give ClO2 or chlorate ion according to Equations 6–7.
Cl2 + ClO2– Ž [Cl2O2] + Cl–
(5)
2 [Cl2O2] Ž 2 ClO2 + Cl2
(6 a)
[Cl2O2] + ClO2– Ž 2 ClO2 + Cl–
(6 b)
[Cl2O2] + H2O Ž ClO3– + Cl– + 2 H+
(7)
Equations 6 a-b are important at high reactant concentrations when the formation of Cl2O2 is
rapid. On the other hand, Equation 7 is more important when the formation of Cl2O2 is slow, such
as at low reactant concentrations or high pH values.
Generators based on the reaction of chlorite ion and chlorine can use either aqueous sodium
chlorite solution or solid sodium chlorite. Chlorine can be in the form of aqueous solution or moist
gas. The reaction between solid sodium chlorite and moist chlorine gas is often used to remove
chlorine from ClO2 gas.
7
Chlorite ion – acid system: Chlorite ion can be protonated to form chlorous acid, which
undergoes a self-decomposition reaction24, 25. The products of this decomposition reaction are ClO2,
chlorate, and chloride ions. This reaction is catalyzed by chloride ion. The stoichiometry of this
reaction changes with the conditions. The two limiting cases are given in Equations 9 and 10. The
actual stoichiometry is given by the linear combination of these equations.
H+ + ClO2– ¾ HClO2
pKa = 1.72 (Ref. 26)
(8)
4 HClO2 Ž 2 ClO2 + ClO3– + Cl– + 2 H+ + H2O
(9)
5 HClO2 Ž 4 ClO2 + Cl– + H+ + 2 H2O
(10)
Equation 9 describes the uncatalyzed reaction and Equation 10 describes the chloride ion
catalyzed decomposition reaction. Because the uncatalyzed reaction produces chloride ion, the
catalyzed pathway becomes significant as the reaction proceeds. A problem with acid-based
generators is that chlorate ion is produced. Another shortcoming of this method is that part of the
chlorite ion is converted to chloride or chlorate ions.
Electrochemical system: These systems present a new technology of ClO2 generators11. Chlorite
ion is oxidized by using electrochemical methods. The overall reaction that takes place in the
generator is the following.
2 ClO2– + 2 H2O Ž 2 ClO2 + 2 OH– + H2
(11)
The advantage of the electrochemical generators is that they require only one chemical for ClO2
generation. These generators are well suited for the generation of low amounts of ClO2, for example
for the generation of ClO2 in the laboratory.
1.2.4. Applications of ClO 2 in water treatment
8
Chlorine dioxide was first used to treat potable water in 1946 at the Niagara Falls, NY water
treatment plant27. The recognition that chlorination of potable water can result in the formation of
toxic chlorinated organic chemicals promoted the use of ClO2. The number of water treatment plants,
which use ClO2 presently, makes up about 5 to 6% of all water treatment plants in the United States4.
Chlorine dioxide is an effective and selective disinfectant. Its application has several advantages
over chlorine. Most importantly, ClO2 does not form halogenated disinfection byproducts and even
reduces the THM formation potential of raw water2, 4. Chlorine dioxide is more effective against some
microorganisms than dissolved chlorine. For example ClO2 can effectively remove Cryptosporidium
cysts that are resistant to chlorine. The disinfection efficacy of ClO2 is independent of pH. However,
above pH 9 the disproportionation of ClO2 becomes significant. This disproportionation reaction
results in a significant decrease in the ClO2 concentration. Thus, to maximize the advantages of the
application of ClO2 as a disinfectant, keeping the pH of the water below 9 is important. Unlike ozone,
ClO2 provides a residual in the water distribution system.
The disadvantages of the application of ClO2 are the following. It is unstable, volatile, and
decomposes when exposed to sunlight. During disinfection, chlorite and chlorate ions form as
inorganic disinfection byproducts2, 11. These byproducts are regulated6, 7 by the D/DBPR. Chlorine
dioxide needs to be generated at the point of application. Due to the need for generators, the cost
of ClO2 can be higher than chlorine4.
Chlorine dioxide can be used for several purposes other than disinfection. It is very effective in
removing manganese(II) and iron(II) from water. Chlorine dioxide is effective in treating taste and
odor problems due to its reaction with phenolic and sulfur containing compounds. Chlorine dioxide
9
can be used for pretreatment of water. Chlorine can be applied subsequently without forming
significant amounts of chlorinated byproducts.
1.2.5. Safety precautions for research laboratories
Chlorine dioxide is toxic28. If ClO2 is inhaled, it can irritate the throat and lungs. The short term
exposure limit for ClO2 is 0.3 ppm (V/V). Because ClO2 is volatile, the concentration of ClO2 above
concentrated solutions can exceed this limit. Therefore, solutions of ClO2 should be prepared and
handled under fume hoods to avoid inhalation.
Sodium chlorite is a strong oxidizing agent. It is important to avoid contact of sodium chlorite
(solid or solution) with combustible materials21. Sodium chlorite generates ClO2 upon contact with
acids and chlorine. For this reason, contact of sodium chlorite with these materials needs to be
avoided. When handling sodium chlorite, the use of protective equipment (eye protection, gloves) is
recommended.
1.3. Research objectives
The main goal of this research was to seek a better understanding of the analytical measurements
of the oxychlorine species ClO2, Cl2O4–, and Cl2O3/Cl2O3–.
Chapters 2 and 3 review some aspects of the measurement of ClO2. In Chapter 2, a newly
developed colorimetric method for the measurement of ClO2 is discussed. The objective of this
Chapter was to determine which commonly encountered interferents present a problem in this new
method and determine the underlying chemical reactions of the interference.
10
In Chapter 3, the formation of the Cl2O4– complex and its effects on the spectrophotometric
measurement of ClO2 are described. The objective was to determine the formation constant and molar
absorptivity of the complex accurately and based on these values provide possibilities for the
correction for the presence of the Cl2O4– complex.
In Chapter 4, a mixed disinfectant solution is discussed. The disinfectant solution is prepared
from dissolved chlorine and ClO2. Increased efficacy is expected of this solution due to the presence
of reactive intermediates that form in the FAC–ClO2 reaction. The objective was to confirm the
increased efficacy of this solution and find out whether the combination of chemical kinetics with
microbiological testing can be used to reduce the number of microbiological tests needed to develop
a new disinfectant.
11
2. Proposed EPA Method 327.0: Determination of ClO 2
and Chlorite Ion in Drinking Water Using Lissamine
Green B and Horseradish Peroxidase (HRP) with
Detection by Visible Spectrophotometry
Any application of ClO2 requires that its concentration be determined accurately. This makes it
important to have reliable analytical methods available. Chlorine dioxide concentrations need to be
determined at various points during its application. For example, it is necessary to measure ClO2 in
the generator effluent to achieve effective use of the generator (low chlorine content, high ClO2
conversion). It needs to be monitored in the treated water to ensure that the required dosage has
been used, but the maximum residual disinfectant level (MRDL), which is set by the US EPA, is not
exceeded. These two examples show that ClO2 concentrations can vary over a wide range. In addition
the interferents present can be very different.
In an effort to provide a reliable, accurate, and easily useable method, the US EPA developed
an analytical method that is able to measure both ClO2 and chlorite ion concentrations. As part of the
promulgation process, a second, independent laboratory is selected to evaluate the performance of
the developed method. This laboratory needs to determine the detection limit, accuracy, and precision
of the method. Furthermore, the comments from this second laboratory are used to improve the new
analytical method. Dr. Gordon’s laboratory was selected to perform this second laboratory evaluation
of the proposed Method 327.0.
In addition to the second laboratory evaluation of the newly developed method, interference
studies were undertaken in conjunction with work at US EPA. The results of the second laboratory
12
evaluation and the interference studies are presented here. Some of the currently existing analytical
methods for ClO2 are reviewed.
2.1. Regulations of ClO2 in potable water
Due to health concerns, the maximum concentration of ClO2 and its inorganic disinfection
byproducts (DBPs), chlorite and chlorate ions, are regulated by the Disinfectants and Disinfection
Byproducts Rule6, 7 (D/DBPR).
The current regulations set the maximum residual disinfectant concentration (MRDL) for ClO2
at 0.8 mg/L. The maximum contaminant concentration (MCL) for chlorite ion, the disinfection
byproduct of ClO2, is 1.0 mg/L. This means that any intended analytical method for the measurement
of these two analytes should have a quantitation level below this concentration, in order to be able
to accurately measure the analyte concentrations.
2.2. Current ClO2 analytical methods
Chlorine dioxide is used as an alternative to chlorine and the first analytical methods for
measuring the concentration of ClO2 were modified chlorine analytical methods. These methods
generally measure the total oxidizing power of the sample and do not have any selectivity for ClO2.
Other analytical methods may suffer from not having the required sensitivity or detection limit to
comply with the current regulations. However, in the case of newly designed methods, it is possible
to keep these objectives in sight and develop a method that satisfies all these requirements.
13
Several good review articles and reports have been published9,
29
on the current analytical
methods. These give a good comparison of the available analytical methods. For completeness and
a better understanding of the newly proposed Method 327.0, a short review of the most important
or most widely used analytical methods is given here.
2.2.1. Ideal Method
The Ideal Method is a theoretical method9. The purpose of this method is to define a generally
accepted set of requirements and based on these to give the possibility of comparing the different
analytical methods. By using the Ideal Method, it is possible to compare different methods objectively
based on their performance and mostly free of personal preferences.
It is important to recognize that in general, none of the analytical methods work equally well in
all samples. Thus, the selection of a suitable method needs to be considered with the parameters of
the sample in mind. For example, if manganese severely interferes with a method, this method can be
the method of choice and work well in water samples in which no manganese is present. On the other
hand, this method is a poor choice for measuring ClO2 in samples that contain high concentrations
of manganese. An example of such samples is the water at Lexington, KY where it is treated with
potassium permanganate.
The requirements defined in this method can be summarized as follows9. The method should
work equally well both in manual and automated mode. It should work for any water sample
regardless of its source. It should have a detection limit of 0.01 mg/L, precision of 0.1%, accuracy
of 0.5%, and a selectivity factor of 500 over generally encountered interferences.
14
2.2.2. Iodometric method
The iodometric method is probably the most widely used analytical method for measuring
oxidizing species in aqueous solutions. In this method, the analyte oxidizes iodide ion to iodine, which
is in turn titrated with standard sodium thiosulfate (Na2S2O3) solution. The titration can be performed
either manually by using starch indicator or by using an automatic titrator.
This method is very useful for measuring the different chlorine species that can be present in
water. Their reactions are described with the following equations:
Cl2 + 2 I– = I2 + 2 Cl–
(pH 7, 2, <0.1)
(12)
2 ClO2 + 2 I– = I2 + 2 ClO2–
(pH 7, 2, <0.1)
(13)
2 ClO2 + 10 I– + 8 H+ = 5 I2 + 2 Cl– + 4 H2O
(pH 2, <0.1)
(14)
ClO2– + 4 I– + 4 H+ = 2 I2 + Cl– + 2 H2O
(pH 2, <0.1)
(15)
ClO3– + 6 I– + 6 H+ = 3 I2 + Cl– + 3 H2O
(pH <0.1)
(16)
From these equations it is clear that by carefully adjusting the pH of the sample when it reacts
with iodide ion, it is possible to distinguish among the various chlorine species9, 30. Depending on the
species that are present, the full method can be modified (simplified) to measure only the species
present. The full method is outlined below. In the experimental section of this chapter, simplified
methods are given for the titration of ClO2, chlorite ion, and chlorine. The experimental details of the
full method are given in Chapter 4.
The samples first react with iodide ion at pH 7. At this pH, chlorine and ClO2 react with iodide
ion. This solution is used in the next step of the titration, when the pH is lowered to 2. At this pH,
chlorite ion and ClO2 oxidize iodide ion. For the next step, a new sample is used with the pH adjusted
to 7. Chlorine dioxide and the volatile portion of chlorine are purged by nitrogen gas from this
15
sample. The sample is titrated to the end point to remove any remaining chlorine. The sample pH is
lowered to 2 and in this case only chlorite ion reacts. If chlorate ion is also present, a new sample is
used, its pH lowered to about 0.1 by using concentrated HCl. At this pH all chlorine species (free
available chlorine, ClO2, chlorite ion, chlorate ion) react with iodide ion.
This is a differential method that can result in significant errors. However, it is a good method
for measuring ClO2 if it is the only species present in the solution. Iodometric titration is used in
laboratories to standardize ClO2 solutions that are used as calibrating solutions for other methods
(e.g., spectrophotometric measurement).
2.2.3. Spectrophotometric method
Chlorine dioxide has a relatively wide absorbance spectrum in the UV/Visible region. The
spectrum is observed in the range31 of about 240-440 nm, having the maximum absorbance12 around
360 nm. This spectrum shows characteristic fine structure due to the 2B1–2A2 electronic transition32.
The molar absorptivity of ClO2 solutions24, 25 at 360 nm is 1250 cm–1M–1. The molar absorptivity of
ClO2 is independent24, 25 of ionic strength (2–4 M), temperature (25–50°C), chloride ion concentration
(up to 0.3 M), and H+ concentration (0.2–4 M).
Several papers have been published where lower values are reported for the molar absorptivity33
of ClO2. The discrepancy can be due to several factors34, including the quality of the photometer.
Many modern spectrophotometers use photodiode array (PDA) detectors, which have limited
resolution (generally about 1 nm). Whereas, the previously cited value (1250 cm–1M–1) has been
determined by using a Cary 14 spectrophotometer, in which the wavelength is selected by a high
resolution monochromator (up to 0.1 nm). The lower wavelength resolution in the modern
spectrophotometers results in averaging of the photometric signal (similar to the moving averaging
16
method). This averaging has a significant influence on the photometric measurement of ClO2, due to
the fine structure of its spectrum.
Figure 2. The change in the measured absorbance of a ClO2
solution as a function of the resolution. [ClO2 ] = 0.0395 M, path
length = 0.0098 cm. — 0.1 nm resolution, — 0.5 nm resolution,
— 1.0 nm resolution, — 2.0 nm resolution
Table 3. Demonstration of the effect of the
resolution on the molar absorptivity
Resolution
A359 nm
Molar abs.
0.1 nm
0.4751
1227.3
0.5 nm
0.4692
1212.1
1.0 nm
0.458
1183.2
2.0 nm
0.4662
1204.3
To illustrate the effect of the resolution of the spectrophotometer on the spectrophotometric
measurement of ClO2, the following measurements were made. The absorbance of the same ClO2
17
solution was measured by using an Olis-Cary 14 spectrophotometer and the resolution of the
photometer was varied between 0.1 nm and 2 nm. The results are shown in Figure 2 and Table 3.
It can be seen from Figure 2, that at lower resolution the absorbance becomes lower in the region
of the spectrum where the fine structure is present. Spectrophotometers do not measure only at the
nominal wavelength, rather in a wavelength interval around the nominal wavelength. Better
resolution means smaller wavelength range. The transmitted light intensity at the nominal wavelength
is the sum of light intensity that is transmitted at all wavelengths in this wavelength range. At lower
resolution this wavelength region is relatively wide, resulting in the change of the measured
absorbance. This averaging effect is insignificant in regions of the spectrum where no sudden changes
occur. However, in the case of ClO2, the region of maximum absorptivity shows a fine structure. At
this wavelength region the averaging effect of the low resolution is significant and can result in lower
absorbance or a shift in the position of the maximum absorbance as compared with high resolution.
These changes can result in molar absorptivity values that vary in a relatively wide range as illustrated
by Table 3. Therefore, it is a good practice to validate the value of the molar absorptivity of ClO2 on
the spectrophotometer that is used.
The spectrophotometric method is relatively free of interferences, easy to perform, and accurate
(if the correct molar absorptivity is used). The typically encountered interferences, e.g., chlorine (free
and combined) or iron do not interfere with this method in concentrations that are generally found
in potable water. However, chlorite ion directly (in high concentration) and indirectly interferes with
this method. The details of these interferences are given in Chapter 3. If the necessary precautions
are taken, this method can be used as a reference method to calibrate other analytical methods. The
concentration range of ClO2 that can be determined in a 1 cm cell at 360 nm is from 4.0×10–5 M
18
(2.7 mg/L) to 1.6×10–3 M (108 mg/L). By using different path length cells (10 cm to 0.1 cm), the
range can be extended to 4.0×10–6 M (0.27 mg/L) and 1.6×10–2 M (1080 mg/L).
2.2.4. Colorimetric methods
These methods are generally based on the reaction between ClO2 and a dye that results in a
decrease in the absorbance of the dye. Due to this reaction, colorimetric methods can be selective for
the measurement of ClO2. The selectivity can be further improved by using gas-diffusion flow
injection analysis (GD FIA)35-37. In addition to the selectivity, the sensitivity of these methods is also
higher than the spectrophotometric measurement, due to the higher molar absorptivity of the dyes.
The working range is limited by the concentration of the dye. The maximum concentration of
ClO2 which can be determined can be increased by increasing the concentration of the dye and
simultaneously decreasing the path length or decreasing the sample size. On the other hand, the
detection limit can be improved by decreasing the dye concentration and simultaneously increasing
the path length or increasing the sample size. However, changing the concentration of the dye and
path length has its practical limits. For this reason, these methods may not be suited for the
determination of very high (e.g., in generator effluents) or very low ClO2 concentrations.
Colorimetric methods can be performed either in manual38-40 or automatic form35-37. Despite the
fact that they can be automated, these methods are not well suited for real-time measurements. This
is due to the generally encountered reaction time, which can be as long as 40 minutes.
A significant problem of the colorimetric methods is the purity of the dyes. The purity of these
dyes can be as low as 40-50% or as high as 90-95%. The impurities present a significant problem
because they also may react with ClO2, thus altering the “stoichiometry” of the ClO2–dye reaction.
19
Furthermore, the stability of the dye either in the solid form or in the reagent can be influenced by the
impurities.
A very good example for the problems that can arise from the impurity of dyes is provided by
the Indigo method for the measurement of ozone41-43. The currently used method43 uses a constant
sensitivity factor to calculate the ozone concentrations. However, recently several papers have been
published that question this practice44, 45. The authors of these papers tested various indigo reagents
and found that in almost all cases, the determined sensitivity factor was significantly different from
the constant value used. This difference is due to the variation in the reagent purity. The purity
changes not just among various sources of indigo reagents, but the purity of the same reagent changes
with time due to decomposition reactions. This variation in the sensitivity factor results in
overestimated ozone concentrations, which translate into increased operational costs and possibly
increased disinfection byproduct formation.
Many colorimetric methods exist for the measurement of ClO2, here only two of them (DPD and
LGB methods) are discussed in detail. Other important methods are only briefly discussed. Table 4
gives an overview of these methods and their parameters.
2.2.5. N, N’-diethyl-p-phenylenediamine (DPD) method
This method originally was developed for the measurement of free and combined chlorine46 and
later adapted for the measurement47-49 of ClO2. The basis of this method is the oxidation of N,N’diethyl-p-phenylenediamine (DPD) to a colored, relatively stable semiquinoid intermediate9. This
intermediate can be further oxidized to a colorless imine. This second step accounts for the fading of
the colored solution.
20
Table 4. Summary of various colorimetric methods used for the measurement of ClO2 .
Reagent
DPD
ACVK
Amaranth
CPR
LGB
lmeasurement (nm)
515/555
548
522
575
616
Detection limit (mg/L)
0.008
0.04
0.006
0.003
0.038
Working range
0.008 to
20 mg/L
0 to
25 mg/L
0.1 to
1.1 mg/L
0.003 to
1.0 mg/L
0 to 0.5 mg/L
Stability of reagent
aqueous reagent
is unstable
NR
3 months
6 months
several
month
% purity of reagent
98a/97b
N/Aa/50b
85a/90b
N/Aa/70b
N/Aa/60b
I
n
t
e
r
f
e
r
e
n
c
e
Free chlorine
%
—
Chloramine
%
not tested
— (up to
17 mg/L)
— (up to
20 mg/L)
—
Chlorite ion
%
—
— (up to
40 mg/L)
— (up to
1000 mg/L)
—
Chlorate ion
NR
—
— (up to
40 mg/L)
— (up to
1000 mg/L)
—
Manganese
%
not tested
—
not tested
—
(MnO2 )
Iron
% (masked by
EDTA)
not tested
—
not tested
—
Other
other oxidizing
species: %
Reference
a
b
ACVK
CPR
DL
DPD
EDTA
LGB
NR
—
%
% (masked by % (masked by
sodium
ammonia
cyclamate)
buffer)
% (masked by
ammonia)
9, 29, 50
Nitrite: — (up
to 1000 mg/L)
29, 51, 52
53
29, 38, 54
Fisher
Aldrich (these purity values are provided only for information and does not represent recommendations)
Acid chrome violet K
Chlorophenol red
Detection limit
N,N’-diethel-p-phenylenediamine
Ethylene diamine tetraacetic acid
Lissamine green B
Not reported
No interference observed. The high end of the tested concentration is given.
Interference is observed.
21
39
Chlorine would interfere with the measurement of ClO2, for this reason glycine is added to the
sample. Glycine reacts with chlorine to form chloroaminoacetic acid, which does not react with DPD.
The pH of the DPD reagent is 6.2–6.5. The absorbance of the resulting solution is measured at
515 nm (or 555 nm) and the ClO2 concentration is determined from a calibration curve. The
calibration curve is created by using standard potassium permanganate solution. Alternatively, the
resulting solution can be titrated with ferrous ammonium sulfate solution until the red color
disappears.
The colored form of DPD has two absorbance maxima at 515 and 555 nm. Traditionally the
peak43 at 515 nm is used to measure the absorbance. In her thesis55, Witte examined the possibility
of using the absorbance change at 555 nm. Her conclusions were that there is no significant difference
in the sensitivity, accuracy, or precision between the measurements at 515 nm or 555 nm.
The interferences of this method include monochloramine, oxidized manganese, chlorite, and
chromate ions. Manganese and chromate ion can be masked by using ethyelenediamine tetraacetic
acid (EDTA).
A significant problem is the fading of the red color of the intermediate. The color change shows
a complex dependence on time. For example, when ozone is added to DPD, the absorbance initially
decreases and then increases41. For this reason it is important that the spectrophotometric
measurements are taken after the same period of time after mixing. Another problem with method
is the stability of the reagent in solution.
The DPD method is a very popular method, and it is used widely in spite of these shortcomings.
This method works well for measuring dissolved chlorine, but the measured ClO2 concentrations are
22
inaccurate50. For this reason, the use of DPD method for the measurement of ClO2 is not
recommended.
2.2.6. Lissamine Green B (LGB) method
This method was originally proposed by Chiswell and O’Halloran39 in 1991. The reagent LGB
was selected due to its high reduction potential. Its reported value is +1.0 V, which is sufficiently high
to eliminate the interference from many commonly encountered interferents. Chiswell and O’Halloran
used cyclic voltammetry to measure the reduction potentials of LGB, ClO2, FAC, chloramine, and
chlorite ion. These measurements were performed under actual working conditions at pH 9.0. The
results indicated that the reduction potential of ClO2 was similar to the reduction potential of LGB
(+0.96 V). The reduction potential of FAC is somewhat lower (+0.76 V) and the other two species
had significantly lower reduction potentials.
The reaction between LGB and FAC or ClO2 was investigated at pH 9.0 in borate buffers.
Chlorine dioxide reacts rapidly with LGB, producing a stable final color. On the other hand, the
reaction between FAC and LGB is much slower, it did not reach completion within an hour. To
eliminate even this low interference from dissolved chlorine, the use of ammonia/ammonium chloride
buffers was suggested. Ammonia reacts with chlorine forming chloramines. Chloramines have lower
redox potential than FAC, thus no interference is expected from chloramines. By using
ammonia/ammonium chloride buffer, mixing 5 mg/L FAC with the LGB solution did not produce
significant absorbance decrease in an hour.
Manganese dioxide and oxidized iron species did not show interference with the measurement
of ClO2. Further advantages of the method can be summarized as follows. The produced color is
stable. This makes it possible to mix the sample with LGB at the sampling point and measure the
23
absorbance in the laboratory. The method is relatively simple, and no pretreatment of the sample is
needed.
Despite its good properties, the LGB method has not been studied further in detail. It has been
used as a reference method53 with other ClO2 analytical methods being compared to the LGB method.
2.2.7. Other colorimetric methods
Amaranth: This method is based on the decolorization of the dye Amaranth. The absorbance
decrease is measured53 at 522 nm. The method was tested in both batch and automated modes. The
response was linear in the 0.1 to 1.0 mg/L concentration range. In batch mode, minimal interference
was observed from chlorite and chlorate ions, monochloramine, and iron(III). Oxidized manganese
shows more significant interference. Aqueous chlorine reacts with Amaranth, but at much slower rate
than ClO2. The use of ammonia/ammonium buffers was tested to eliminate the interference of FAC.
The results indicated that this buffer is effective in masking FAC.
In addition, gas diffusion flow injection analysis was tested to eliminate the interference of FAC.
The results showed that the use of GD-FIA makes the Amaranth method selective for measuring ClO2
in the presence of FAC. The determined selectivity factor was on the order of 1000. The method
shows good promise: it has good selectivity and a low detection limit.
Acid chrome violet potassium salt (ACVK): ACVK is decolorized by ClO251, 52, 56. No interference
is observed from free or combined chlorine, chlorite and chlorate ions. Ozone does react57, 58 with
ACVK. However, this is not a significant problem because ClO2 and ozone react with each other59.
The problems with this method include the complicated reagent preparation and possible problems
with the stability of the reagent.
24
Clorophenol red (CPR): This method was first proposed by Wheeler et al 40 in 1978. Since then
several improvements were made to this method38, 54, 60. It is a relatively selective method for ClO2.
Chlorine, chlorite, iron, and manganese do not interfere.
2.2.8. Electrochemical methods
The use of a voltammetric rotating membrane electrode has been reported61. Here, the membrane
provided the selectivity of the electrode and no interference was observed from hypochlorite, chlorite,
chlorate, and permanganate ions. No detection limit has been reported, but the authors were able to
measure ClO2 concentration around 0.30 mg/L. The electrode gave similar results to the chlorophenol
red method. The 90% response time was about 1 minute.
Glassy carbon and platinum amperometric sensors were used for the measurement62 of ClO2.
These sensors showed a good response for ClO2 and chlorite ion. However, they were tested only
at pH 4 in the water used for pulp bleaching, and no detailed parameters were given.
Quentel et al described an interesting electrochemical method for the measurement of ClO2 at
low levels. In separate experiments two dyes, Alizarin Red S63 and Indigo-Carmin64, were used to
determine the ClO2 concentration. The dyes reacted with ClO2 that resulted in a decrease in their
concentration. At high ClO2 concentrations, this decrease can be measured by spectrophotometric
measurement. At low ClO2 concentrations (few :g/L), the concentration decrease of the dye is
measured by voltammetry. To improve the sensitivity of the method at low concentrations, the dyes
were electrochemically preconcentrated on a mercury drop electrode. The selectivity of the method
is similar to the original colorimetric methods. Chlorite ion and free or combined chlorine do not
interfere under the experimental conditions.
25
2.3. Experimental
2.3.1. Reagent water
All reagent solutions were prepared by using triple-distilled water (TDW). Water from a
Barnstead Nanopure system with at least 18.4 MS cm-1 resistance was doubly-distilled in an all-glass
Barnstead Fi-Stream still. TDW was stored in Nalgene carboys.
2.3.2. Generation of ClO 2
The various ClO2 generation methods have been overviewed in Chapter 1. Here, only the
laboratory preparation is detailed. Chlorine dioxide was generated by the oxidation of chlorite ion
with persulfate ion according to Equation 17.
2 ClO2– + S2O82– Ž 2 ClO2 + 2 SO42–
(17)
Chlorine dioxide solutions were generated based on a previously published method65. Fifty mL
of a 16% (m/m) sodium chlorite solution was mixed with 100 mL of a 4% (m/m) potassium persulfate
solution in a gas washer. The ClO2 formed was purged from the solution by using pre-purified
nitrogen and absorbed in chilled TDW.
All ClO2 solutions were transferred to amber bottles with Teflon lined caps. The bottles were
filled so that there was no headspace to avoid evaporation and decomposition of ClO2. The solutions
were stored in a refrigerator below 6°C. Prior to use, the ClO2 solutions were transferred to a
shrinking bottle and titrated by using the iodometric procedure.
26
2.3.3. Carbonate free sodium hydroxide solutions
Even the highest purity solid sodium hydroxide is contaminated with sodium carbonate due to
adsorption of CO2 from the air. In general, the presence of sodium carbonate in sodium hydroxide
solutions is undesirable. Carbonate free sodium hydroxide solutions were prepared from a 50%
NaOH solution66 in which sodium carbonate is highly insoluble. Calculated volumes of the
50% NaOH solution were diluted to the required volume by using CO2 free TDW. These solutions
were used within two days of preparation.
2.3.4. Preparation of dissolved chlorine solutions
Chlorine solutions were prepared by bubbling chlorine gas through a 0.1 M carbonate free NaOH
solution, which was freshly prepared as described above. The pH of this FAC solution was set to 11
to minimize the decomposition8 of FAC. The solutions were stored in Nalgene bottles below 6°C. The
concentration of FAC was determined by iodometric titration.
2.3.5. Preparation of monochloramine
Chloramines are formed in the reaction of ammonia with chlorine in aqueous solution. The
formation of chloramines is described with the following simplified equations9.
NH3 + HOCl Ž NH2Cl + H2O
(18.a)
NH2Cl + HOCl Ž NHCl2 + H2O
(18.b)
NHCl2 + HOCl Ž NCl3 + H2O
(18.c)
The distribution of these species is controlled by several factors, including temperature, pH, and
the ammonia to chlorine ratio. At high ammonia to chlorine ratios mainly monochloramine is formed.
With increasing chlorine concentration (decreasing ammonia to chlorine ratio) monochloramine can
27
react further with chlorine to form dichloramine. Dichloramine is not stable, but the detailed
mechanism of its decomposition is not known9. Upon further addition of chlorine to the solution only
a small amount of nitrogen trichloride is formed.
Based on this information, it is necessary to have excess ammonia in order to form only
monochloramine. The optimum pH for the formation of monochloramine is 8.0–8.5. Monochloramine
solutions were prepared9 by mixing dilute aqueous chlorine solution with ammonium chloride
solution. The concentration of the ammonium chloride solution was selected so that it contained at
least three times more ammonium ion than the amount of chlorine added in molar units. The pH of
the ammonium chloride solution was adjusted to 8.3 with phosphate buffer before mixing with the
chlorine solution. The dilute chlorine solution was added dropwise to the ammonium chloride
solution, while stirring vigorously. The monochloramine solutions were stored in plastic bottles below
6°C. Fresh solutions were prepared weekly. The concentration of this solution was determined by
iodometric titration.
2.3.6. Titration of chlorine and monochloramine solutions
Because these solutions contained only FAC or monochloramine, but not the combination of
them, it was possible to determine their concentrations in one step. The samples were added to 30 mL
of 0.1 M pH 7.0 phosphate buffer. To this solution, about 0.5 g solid KI was added and the formed
iodine was titrated with standardized 0.1 M sodium thiosulfate solution.
2.3.7. Titration of ClO 2 and chlorite ion solutions
The full iodometric procedure was simplified because only a single analyte was present in these
solutions. Chlorine dioxide or chlorite ion samples were added to about 30 mL water, which was
28
acidified by using 2.0 mL of a 2.5 M HCl solution. About 0.5 g KI was added to the solution, and
the iodine formed was titrated with standard sodium thiosulfate solution.
2.3.8. Preparation and titration of potassium permanganate solutions
About 0.15 g of potassium permanganate was dissolved in one liter TDW. The solution was
heated to boiling and kept hot for about one hour. It was covered and allowed to stand overnight.
The solution was filtered on the next day by using a fine porosity sintered glass filter. The final
solution was stored in an amber bottle.
The permanganate solution was standardized with sodium oxalate67. Sodium oxalate was
dissolved in water and 3 M H2SO4 was added. This solution was heated to about 80 to 90°C and
titrated with the KMnO4 solution. The potassium permanganate solution was added slowly, making
sure that all of it reacted before adding the next increment. This is necessary due to the relatively slow
reaction between oxalate and permanganate ions. It was necessary to keep the temperature of the
solution above 60°C throughout the titration.
2.3.9. Other reagents
The LGB, horseradish peroxidase solutions, and the glycine/citric acid buffer were prepared as
described in the proposed68 Method 327.0. Chlorite ion standards were prepared by dilution from
Absolute Standards, # 54109 sodium chlorite standard. The concentration of this standard was 1 g/L.
2.3.10. Shrinking bottle
A shrinking bottle69 is a modified syringe that can be used for the delivery of accurate volumes
of solutions containing volatile compounds. Figure 3 shows a schematic diagram of the shrinking
bottle. The shrinking bottle allows the delivery of the solution without creating a headspace, thus
29
preventing evaporation loss. It is made from a 50 mL syringe. The plunger of the syringe is attached
to a precision screw that allows the accurate delivery of known volumes. It can be calibrated by
weighing the water that is delivered by one full turn of the screw. Less than 1% ClO2 was lost daily
when the solutions were stored in the shrinking bottle.
Figure 3. Drawing of a shrinking bottle69. A precision screw, B - brass frame, C - retaining
springs, D - guide for screw, E - 50 mL
syringe
2.3.11. Other equipment
All titrations were performed by using a Radiometer ABU 93 Triburette station, which was
controlled by a Radiometer VIT 90 Videotitrator unit. Radiometer M21Pt platinum and Radiometer
Ref 401 calomel reference electrodes were used to follow the titrations. The pH of the solutions was
checked by an Accumet glass electrode connected to the VIT 90 Videotitrator unit.
30
Spectrophotometric measurements were taken on an Agilent 8453 spectrophotometer. The
spectrophotometer is equipped with a photodiode array detector with 1 nm resolution. Non-volatile
and non-corrosive solutions were transferred by a Rainin EDP Plus electronic pipette. The pipette
was calibrated regularly by gravimetric procedure.
2.4. The proposed LGB method
As part of a major project to provide water utilities with the necessary analytical methods for
compliance with Stage 2 Disinfectants and Disinfection Byproducts Rule6 (D/DBPR), US EPA has
developed a new method68 for the measurement of ClO2 and chlorite ion. Some details of the method
and a short description of the procedure are provided here.
The purpose of this method is to provide water utilities with a simple, accurate method for the
measurement of ClO2 and chlorite ion. The emphasis was to develop a method which can be used
under field conditions. Colorimetric methods qualify well for this purpose due to their (typically)
simple procedure and the availability of pocket colorimeters and field photometers.
Initially four colorimetric methods were considered, ACVK, Amaranth, CPR, and LGB. From
these methods, LGB was selected due the low number of interferences and the highest sensitivity39, 70.
The maximum absorbance of this dye is dependent on the pH, but it is in the 600 nm region at all pH
values. In this region of the visible spectrum usually less interference is observed than at shorter
wavelengths. For example, permanganate ion, a typical colored interference, does not interfere. The
structure of LGB is shown in Figure 4.
31
Figure 4. The chemical structure of LGB
Chlorine dioxide reacts with LGB, but no reaction is observed between LGB and chlorite ion.
However, chlorite ion can be oxidized to ClO2 by Horseradish Peroxidase (HRP)71, 72. The enzyme
shows maximum activity in the pH 6.0–6.5 region. For this reason it was necessary to modify the
original LGB procedure39, which used ammonia/ammonium chloride buffers at pH 9.0 to eliminate
the interference of dissolved chlorine. The current method uses glycine to eliminate the interference
of FAC and the pH is adjusted by using a citric acid buffer. The ClO2 and chlorite ion concentrations
are determined by using the absorbance decrease at 633 nm.
Calibration curves are constructed by using chlorite ion standards. The reason is that ClO2 is
volatile and reactive, thus it is not suited for the preparation of reliable standard solutions, especially
under field conditions. Furthermore, the good correlation between the chlorite ion concentration and
the ClO2 concentration formed allows the construction of calibration curves for ClO2 based on
chlorite ion standards.
The final procedure of the proposed Method 327.0 uses amber vials as volumetric glassware for
mixing of the samples with the LGB/HRP reagent. The vials are calibrated by a gravimetric
procedure.
32
Relatively rigorous quality control procedures are used in this method. A Continuing Calibration
Check Standard (CCCS) is analyzed in each batch of samples (about 10 samples). If the percent
recoveries for these Continuing Calibration Check Standards are not in the range of 70 to 130% (or
50 to 150% for the 0.25 mg/L standard) as established by the US EPA, the samples in the batch need
to be reanalyzed.
The procedure is outlined below. The sample is transferred to an amber vial and an aliquot is
removed. The same volume of pH 6.0–6.5 citric acid/glycine buffer is added to the sample. Following
the mixing of the solution, another aliquot is removed and the combined LGB/HRP solution is added
to the sample and mixed. After 20 to 40 minute reaction time, the absorbance of the solution is
measured. The concentrations of ClO2 and chlorite ion are calculated from the absorbance difference
between the samples and blank measurements. To determine only the chlorite ion concentration, the
sample is purged with nitrogen (or other inert gas) before transferring into the amber vials. This step
removes ClO2. Following this, the sample is analyzed as described above. The absorbance reading of
this step gives the chlorite ion concentration, which is subtracted from the ClO2 plus chlorite ion
concentration to calculate the ClO2 concentration.
2.5. Results of the second laboratory experiments
In order to assess the properties of the newly developed method, US EPA selected our
laboratory to perform a second laboratory evaluation of the newly developed method. The purpose
of the second laboratory experiments is to confirm the performance of the proposed analytical
method, objectively assess its properties and help the US EPA to improve the method. The results
of the second laboratory experiments are summarized in the following sections.
33
A brief description of the experimental procedure is given in the previous section. More details
can be found in the proposed method68, which is available from US EPA. During this work, however,
more standard solutions were included than in a real world application. The reason for this was to
have a more stringent quality control of the experiments.
The analysis of the samples for the initial assessment of the performance (detection limit,
accuracy, and precision) of the method was performed on five consecutive days. On three days,
laboratory reagent water (TDW) was used and on the other two days Oxford tap water. The purpose
of using tap water was to demonstrate the properties of the method in a real water matrix. The
Oxford tap water is chlorinated and no ClO2 is used during its treatment. Thus, no ClO2 or chlorite
ion is present in the tap water. This was confirmed by comparing the absorbance of LGB solutions,
which were added to tap water, with the absorbance of LGB solutions that were mixed with TDW.
The measured absorbances were not significantly different as compared by t-test at the 95%
confidence level.
A calibration curve was measured each day by using five calibration standard solutions. The
chlorite ion concentrations of these samples were 0.25 mg/L, 0.51 mg/L, 1.01 mg/L, 1.52 mg/L, and
2.03 mg/L. The measured parameters of the calibration curves are summarized in Table 5.
Table 5. Parameters of the calibration curves on various days.
Day 1
Reagent Water
Day 2
Reagent Water
Day 3
Reagent Water
Day 4
Tap Water
Day 5
Tap Water
Slope
0.391
0.365
0.398
0.396
0.371
Intercept
-0.008
-0.004
-0.021
-0.039
-0.002
R2
0.996
0.988
0.995
0.993
0.998
34
During the measurements, Continuing Calibration Check Standard solutions were run to check
that the calibration was still valid. These standards included both chlorite ion and ClO2 standards (in
separate solutions). However, only chlorite ion standards were used to check the validity of the
calibration.
The absorbance change at 633 nm was compared for the five calibration curves (obtained on
different days). The five calibration curves were compared in pairs by using multiple linear model
regression. For each pair of lines, the tentative model contained three independent variables: the ClO2
concentration, an indicator variable for one of set of measurements (reference line), and the product
of the previous two. Each of these models assumes that the variances of the two corresponding error
terms are the same, but permits the two corresponding regression lines to have different slopes and
intercepts.
The results of the fitting are shown in Table 6. A fit of each of these models to the data (using
the stepwise procedure of SPSS73) showed the following in each model. First, the only significant
variable is the ClO2 concentration at the 95% confidence level. Consequently, neither the indicator
variable, nor the product of the indicator variable and the ClO2 concentration are significant at the
95% confidence level. This indicates that the two regression lines compared are not statistically
different from each other. Also note, that the constant (intercept of the lines) is not significant in any
of these models (p>0.05 in each model), indicating that each of the regression lines does go through
the origin.
These results show that the calibration curves are not statistically different. Thus, all calibration
data can be fitted in one linear model. The equation of this overall calibration curve is
DA633 nm = 0.384×[ClO2–]
35
(R2 = 0.996)
(19)
Table 6. Comparison of the calibration curves for the LGB method by
using multiple linear model regression.
ClO2
Intercept
Calibration curves
compared*
Slope
p
Value
p
1
2&1
0.378
<0.001
-0.006
0.743
0.989
2
3&1
0.394
<0.001
-0.014
0.278
0.995
3
3&2
0.382
<0.001
-0.012
0.505
0.989
4
4&1
0.393
<0.001
-0.023
0.176
0.992
5
4& 2
0.380
<0.001
-0.022
0.244
0.989
6
4&3
0.397
<0.001
-0.030
0.084
0.993
7
5 &1
0.381
<0.001
-0.003
0.806
0.995
8
5&2
0.368
<0.001
-0.001
0.928
0.992
9
5&3
0.385
<0.001
-0.009
0.460
0.995
10
5&4
0.384
<0.001
-0.019
0.211
0.993
Model #
R2
* The second data set is the reference.
This result shows that the daily calibration of the method may not be necessary. The method probably
could be modified to create a calibration curve only when the combined LGB–HRP reagent is
prepared and Continuing Calibration Check Standards are measured regularly to ensure the accuracy
of the measurements. If deviation is observed in the recovery of the CCCS, a new calibration curve
is created. If no deviation is observed, the initially created calibration curve could be used for two
weeks (the shelf-life of the reagent). This could significantly simplify the method.
However, further studies are needed to confirm this assumption because the current results are
based on a limited data set. Comparing the calibration on different days for more than one LGB
reagent would also be necessary. In addition, the LGB–HRP solutions were prepared freshly every
week. Thus, no conclusions can be drawn on the stability of the reagent for the period of two weeks.
36
2.5.1. Method detection limit
The detection limit of the new method was determined both for ClO2 and chlorite ion by
analyzing seven replicates of samples at 0.25 mg/L. The detection limits were determined for cases
when only one of the analytes was present and when both analytes were present. The detection limit
was calculated according74 to Equation 20.
DL = s×t(n-1,1-" = 0.99)
where s
(20)
standard deviation of the replicate measurements
t(n-1,1-" = 0.99)
Student’s t value for the 99% confidence level with n-1 degrees of freedom
n
number of replicates
The detection limits for the two species are shown in Table 7. The detection limit for ClO2 and
chlorite ion are at about the same level of 0.1 mg/L. These detection limits are about an order of
magnitude higher than the detection limit suggested by the ideal method. However, this detection
limit is sufficiently low to be able to measure ClO2 and chlorite ion at their respective regulatory
concentrations.
Table 7. Detection limits for chlorite ion and ClO2 in reagent water.
Detection limit (mg/L)
[ClO2–] (mg/L)
[ClO2] (mg/L)
0.25
ClO2–
ClO2
N/A
0.1
N/A
0.25
0.8
0.19
N/A*
N/A
0.25
N/A
0.16
1
0.25
N/A*
0.21
* Detection limit was calculated only at analyte concentrations that are close to the
expected detection limit.
37
The detection limit of both species is influenced by the presence of the other analyte. A possible
reason for this increase can be due to the fact that this method is a differential method that can result
in higher errors when more than one species is determined.
2.5.2. Recoveries of the samples
Figures 5 and 6 show the distribution of the percent recoveries of the chlorite ion and ClO2
samples. These include the previously mentioned Continuing Calibration Check Standards and
samples that contained only a single analyte. The symbols represent the measured values, the bars
represent one standard deviation and are centered on the average value. Percent recovery is defined
by the following equation.
(21)
Figure 5 shows that the distribution of the percent recoveries for the chlorite ion samples only
slightly changes with the fortification level. At 0.25 mg/L, the standard deviation is relatively high,
but at 2.0 mg/L, the standard deviation is lower. The figure shows that the mean percent recovery
is close to 100%, indicating no problem with the accuracy of the method.
38
Figure 5. The percent recoveries of pure chlorite ion standards, which
were determined during the second laboratory testing. — 0.25 mg/L
chlorite ion,  1.0 mg/L chlorite ion, • 2.0 mg/L chlorite ion.
Figure 6. The percent recoveries of pure ClO2 standards, which were
determined during the second laboratory testing. — 0.25 mg/L ClO2,
 0.8 mg/L ClO2, • 2.0 mg/L ClO2.
39
Figure 6 shows that the distribution of the percent recoveries of the ClO2 samples is greatly
dependent on the fortification level. The standard deviation at 0.25 mg/L is significantly higher than
at 2.0 mg/L. Furthermore, the figure shows that the mean recovery is different from 100%, especially
at 0.8 mg/L and 2.0 mg/L. This indicates a possible problem with the accuracy of the method for the
measurement of ClO2.
A possible reason for the consistently higher ClO2 concentration is a problem with the calibration
curve. The calibration curves were constructed by using sodium chlorite standards. Chlorite ion is
converted to ClO2 by HRP.
(22)
Here " is a conversion factor that has the maximum value of 1. From Equation 22, the equation
of the calibration curve is
(23)
where slope1
slope2
the slope of the calibration curve for ClO2
the slope of the calibration curve for chlorite ion
If " equals one, the slopes of the ClO2 and chlorite ion calibration curves are the same. Studies,
conducted by US EPA, showed that the conversion factor under the used conditions is unity.
However, in other studies of the HRP-chlorite ion reaction, lower than unity conversion factors have
been reported72. If " is lower than one, the slope of the ClO2 calibration curve is higher than the slope
of the chlorite ion calibration curve. Thus, the determined ClO2 concentrations would be higher than
the true value. If " is constant in the range of the calibration curve, the determined ClO2
concentrations can easily be corrected by multiplying them by ".
40
Based on the current results it appears that the conversion factor (") is lower than unity.
However, to confirm this possibility, further studies are needed. It is important to determine the
stoichiometry of Equation 22 and determine if this stoichiometry is not changing with chlorite ion
concentration.
2.6. Interference studies
An interference study was performed jointly with the US EPA. The results of this interference
study, which were obtained in our laboratory, are presented here.
Throughout this section the following conventions are used. Chlorine dioxide concentration that
is achieved by spiking TDW with concentrated ClO2 stock solution is called the true ClO2
concentration. At some places it is also called the added ClO2 concentration. The concentration of
the ClO2 stock solution was determined by iodometric titration as described in the Experimental
section of this Chapter. The determined ClO2 concentration is the ClO2 concentration that is
calculated from the absorbance decrease of the LGB solution at 633 nm.
2.6.1. Interference in analytical measurements
Even though interference studies are routinely performed for analytical methods, interference is
not a well-defined, universal term. The recommendations of IUPAC are summarized below75.
The main problem with the general definition of interference is that it should apply to both
qualitative and quantitative analysis. The following definition of interference is given in the IUPAC
recommendation: “An interfering substance for an analytical procedure is one that causes a
predeterminate systematic error in the analytical results.”
41
The authors caution that “the allowable magnitude of the systematic error should be fixed
beforehand.” This practice makes it possible to make objective judgments of the performance of the
analytical method. In this aspect, it is similar9 to the “Ideal Method” in that it gives the possibility of
unbiased decisions based on the true performance of the analytical method, rather than on personal
preferences.
The concentration of the interferent, the presence of other compounds in the sample, and the
concentration of the analyte affect whether a compound interferes or not. The magnitude of the
interference in many cases is not linearly dependent on the concentration of the interferent 76 . The
presence of other compounds in the sample can significantly affect the interference (similar to the
matrix effect).
Wilson76 recommends “that the effect of any substance should be estimated for at least two
concentrations” of the analyte. The two recommended concentrations are the lowest and highest
concentrations of the working range of the analytical method. If the results at these two
concentrations show a difference in the interference, further testing is necessary at intermediate
analyte concentrations.
Testing for interference is recommended at a minimum of one interferent concentration that is
higher than the expected maximum concentration of the interferent in the samples. For ClO2 analytical
methods, which are used in potable water, the maximum expected interferent concentrations can be
estimated by the corresponding federal regulations (e.g., Safe Drinking Water Act5, or Disinfectants
and Disinfection Byproducts Rule6, 7). However, in order to gain more information about the
interference, it is necessary to measure more than one interferent concentration.
42
For the purpose of this Chapter, interference is defined as an absorbance change at 633 nm (the
position of the maximum absorbance of LGB) that is greater than three times the standard deviation
of blank solutions (±0.01×3 absorbance units). The measured absorbance in the presence of the
interferent is compared with the measured absorbance of a solution with the same composition, but
without the interferent. The various sources of the interference can be summarized as follows.
!
Changes in the free reagent concentration: this includes the oxidation of the LGB similar
to the ClO2 oxidation. The other possibility is a complex formation with LGB. Both
processes decrease the free LGB concentration resulting in a decrease in the absorbance at
633 nm. Thus, this is always a positive error.
!
Other interferences include colored interferents which significantly absorb light at 633 nm.
This type of interference always increases the absorbance, resulting in negative error in the
ClO2 concentration.
The magnitude and the sign of the observed interference are the sum of each of these processes.
It is possible for a certain concentration of the interferent that the errors due to each process may
cancel, resulting in a determined ClO2 concentration that is not significantly different from the true
ClO2 concentration. This can be easily identified if the determined ClO2 concentrations are compared
with the true ClO2 concentration at a range of interferent concentrations, instead of making this
comparison at each interferent concentration separately.
2.6.2. Demand vs. interference
There are compounds in raw water that readily react with ClO2, e.g., phenolic compounds, sulfite
ion. These compounds consume ClO2, resulting in a ClO2 demand. Under certain conditions, demand
can be incorrectly interpreted as interference. An explanation for this is given below.
43
Consider a raw water that contains phenolic compounds. If this water is treated with ClO2, the
reaction of the phenolic compounds with ClO2 results in a demand for ClO2. In addition, it is assumed
that no interference process takes place, only demand is present. If ClO2 is applied in excess, some
residual ClO2 concentration will remain in the treated water. This residual concentration, however,
is lower than the originally applied ClO2 concentration.
Here, the true and added ClO2 concentrations are different. The added ClO2 concentration is the
ClO2 concentration that was originally added to the raw water and could be measured if phenolic
compounds were not present in the water. The true ClO2 concentration is the ClO2 concentration,
which is present in the treated water and could be measured with an ideal method with which none
of the components of the treated water interfere. The measured ClO2 concentration is the ClO2
concentration that is determined from the absorbance of the LGB solution at 633 nm.
If the measured ClO2 concentration is compared with the added ClO2 concentration, the
conclusion would be that phenolic compounds interfere with the measurement. The reason for this
is the fact that the measured ClO2 concentration at most can be equal to the true ClO2 concentration,
which in this case is lower than the added ClO2 concentration due to the ClO2 demand of the phenolic
compounds. Here, the ClO2 demand is incorrectly interpreted as interference.
However, if the measured ClO2 concentration is compared with the true ClO2 concentration, the
conclusion would be that phenolic compounds do not interfere. This is due to the assumption that no
other interference process occurs and the true and measured ClO2 concentrations are not different.
For this reason, it is important to consider the possible reactions between ClO2 and the potential
interferent.
44
It is possible that some strong oxidizing compounds react with chlorite ion and chlorine dioxide
is formed from this oxidation reaction. These reactions decrease the observed demand for ClO2 by
recycling chlorite ion into ClO2. Thus, the observed ClO2 demand is the sum of all these processes.
It may not be possible to determine the contribution of the individual reactions to the demand.
However, these contributions can be estimated based on the known reactions between the interferent
and ClO2 or chlorite ion. The important distinction is that demand and interference are quite different
chemical processes.
2.6.3. The interferences studied
The effect of free available chlorine (FAC), chloramine, chlorate (ClO3–), iron(II), manganese(II),
and manganese(VII) (permanganate) ions were studied on the measurement of ClO2. For this reason
the LGB reagent was slightly modified. The buffered LGB solution was mixed with a diluted buffer
solution, which was prepared following the same procedure as for the HRP solution. However, no
HRP was added, because no chlorite ion concentrations were measured.
Based on the previous recommendations, the testing procedure was devised as follows. All
interferents were measured at four different concentrations. At least one concentration was above the
maximum contaminant level (MCL) or maximum residual disinfectant level (MRDL) of the
interferents. At each interferent concentration three different measurements were taken: only
interferent, interferent + 0.25 mg/L ClO2, and interferent + 2.0 mg/L ClO2. All solutions were
measured in triplicate.
The determined ClO2 concentrations in the presence of the interferents were compared in each
case with the determined ClO2 concentrations in the absence of the interferent and with the true ClO2
45
concentration. The comparison was made by using the t-test at the 95% confidence level. Any
possible reactions between ClO2 and the interferents are described.
2.6.4. Chlorate ion interference
The measured range for chlorate ion was from 1 to 10 mg/L. In these experiments, sodium
chlorate was used. The results of the measurements are summarized in Figure 7 and Table 8
The results show that when only chlorate ion is present in the solution, the measured absorbances
are not significantly different from the blank solutions, and no trend is observed. The determined ClO2
concentrations from these absorbances are well below the detection limit of the method. This means
that chlorate ion does not react with LGB at the concentration range tested.
Figure 7. The change of the determined ClO2 concentration
with chlorate ion concentration.  no ClO2 added, Ž
0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added. The dashed
lines show 0, 0.25, and 2.0 mg/L determined ClO2
concentrations.
46
Table 8. Determined ClO2 concentrations in the presence of chlorate ion, in mg/L units.
[ClO3–] (mg/L)
Added ClO2 (mg/L)
0
0.25
2
0
N/A
0.26 (±0.04)
1.88 (±0.05)
0.92
-0.01 (±0.04)
0.32 (±0.03)
1.96 (±0.01)
1.83
-0.09 (±0.14)
0.32 (±0.03)
1.97 (±0.03)
4.59
0.04 (±0.01)
0.28 (±0.02)
1.92 (±0.034
9.18
0.03 (±0.04)
0.21 (±0.03)
1.86 (±0.03)
average*
0.04
0.28
1.93
corrected**
N/A
0.24
1.89
± represents the standard deviation of three samples
* in the case of blank solutions (no ClO2 added, second column) it is the standard error from zero. This
is calculated by averaging the absolute values of the determined ClO2 concentrations.
**calculated by subtracting the average determined ClO2 concentration for blank solution from the average
determined ClO2 concentration
In the case when ClO2 and chlorate ion are present in the solution, the determined ClO2
concentrations were compared with the determined ClO2 concentration at the same level of ClO2
added, but no chlorate ion was present. The mean ClO2 concentrations were compared by using t-test.
The results show that there is no significant difference between the determined ClO2 concentrations
(at the 95% confidence level) in the presence and absence of chlorate ion. In the case of 0.25 mg/L
added ClO2 concentration, the average of the determined ClO2 concentrations is very close to the true
ClO2 concentration. The accuracy of the measurement of 0.25 mg/L ClO2 is further improved by
correcting by the average determined ClO2 concentration for the blank solution, as the last row of the
Table shows.
In the chlorate ion range tested, a small decrease is observed in the determined ClO2
concentration at 0.25 mg/L added ClO2 concentration. This decrease can not be assigned to any of
47
the previously described interference processes for the following reasons. Chlorate ion does not
absorb light in the visible region, thus it does not interfere by increasing the absorbance at 633 nm.
Furthermore, the results in the absence of ClO2 indicate that chlorate ion does not react with LGB.
Thus, this interference mechanism can be excluded. Finally, chlorate ion is not expected to react with
ClO2. Thus, the decrease in the determined ClO2 concentration can not be explained by the demand
of chlorate ion for ClO2.
At 2.0 mg/L added ClO2, the determined ClO2 concentrations in the presence of chlorate ion are
not significantly different from the determined concentration of the blank solution. The results are in
good agreement with the true ClO2 concentration. Even though, in this case correcting for the blank
makes the accuracy lower, the percent recovery is about 95%. This recovery still can be considered
good. Similar to the results when 0.25 mg/L ClO2 was added, the determined ClO2 concentrations
are decreasing at 2.0 mg/L added ClO2 concentration.
Conclusions: Chlorate ion does not show interference with the measurement of ClO2. No
interference was observed neither in the absence nor in the presence of ClO2. Based on these results,
chlorate ion does not appear to react with LGB or ClO2 in the tested range.
2.6.5. Iron(II) interference
The secondary drinking water standard for iron77 is 0.3 mg/L. However, this standard is not an
enforceable level and only serves as a guideline to avoid aesthetic (color or odor) problems. For the
interference measurement with iron(II), ferrous ammonium sulfate solutions were used in the 1.0 to
10.0 mg/L concentration range.
In water, iron(II) can be oxidized to iron(III). This oxidation is not always complete, resulting
in mixtures of iron(II) and iron(III). In the current measurements a similar situation can be
48
encountered. Thus, it is important to consider the possible interference from both iron(II) and
iron(III). The possible reactions of iron(II) and iron(III) with the sample and reagent components are
summarized below.
!
Reaction with LGB
"
Complex formation
"
Redox reaction
!
Complex formation with citrate and hydrogen citrate ions
!
Reaction with ClO2 (demand for ClO2)
!
Reaction with chlorite ion: formation of ClO2, decreases demand
Figure 8 shows the measured spectra of the LGB solutions at various ClO2 concentrations.
Figure 9 shows the measured spectra of the LGB solutions at various Fe(II) concentrations.
Figure 8. Measured spectra of the LGB solution after the addition of ClO2
solutions. The inset shows the spectral region of 275 nm to 325 nm. — No
ClO2, — 0.5 mg/L ClO2, — 1.0 mg/L ClO2, — 2.0 mg/L ClO2
49
Figure 9. Measured spectra of the LGB solution after the addition of
iron(II) solution in the absence of ClO2 . The inset shows the spectral
region 275 nm to 325 nm. — No Fe(II), — 1.00 mg/L Fe(II), —
2.00 mg/L Fe(II), — 5.00 mg/L Fe(II), — 9.99 mg/L Fe(II)
Comparison of Figures 8 and 9 shows that the absorbance below 400 nm increases with the
addition of increasing amounts of iron(II). This absorbance increase is in contrast with the absorbance
decrease when ClO2 is added to the LGB (Figure 8). Another difference between the measured
spectra in the two cases is the isosbestic point around 280 nm. The isosbestic point is present in the
measured spectra when only ClO2 is added to LGB, but it is absent from the spectra of LGB mixed
with iron. This difference is very likely associated with the increased absorbance upon the addition
of iron(II) solutions.
The absorbance increase below 400 nm can be the indication of the formation of a new species.
This new species is possibly an iron containing complex. This can be confirmed by plotting the
absorbance at 303 nm (the maximum of one of the peaks in this region) as a function of iron(II)
concentration. The plot is shown in Figure 10.
50
Figure 10. The absorbance change at 303 nm as the function
of Fe(II) concentration. The line is the least squares fitted line.
The equation of this line: Abs303 nm = 0.0310×[Fe2+] + 0.255,
R2 = 0.999
The linear relationship between the absorbance at this wavelength and the iron concentration
indicates that iron is a component of this complex. The other component of this complex can be the
hydrogen citrate/citrate ion, glycine, or LGB. If the dye forms the complex with iron, that can be a
significant interference due to the decrease in the “free” LGB concentration. To determine if the
complex is formed with LGB, iron solutions were mixed with LGB solutions that did not contain the
citric acid/glycine buffer. The same procedure was followed as described in the proposed Method
327.0, but instead of the buffer solution, TDW was added. The measured spectra are shown in
Figure 11.
In this case no absorbance increase was observed below 400 nm. Thus the results indicate that
the previously measured absorbance increase below 400 nm is not due to an iron(II)–LGB complex.
When no citric acid buffer was present, the absorbance of the peak at 303 nm decreases with
increasing iron concentration. In addition, the absorbance decreases at 633 nm. The absorbance
51
Figure 11. Measured spectra of the LGB solution after the addition of
Fe(II) solution in the absence of glycine/citric acid buffer and ClO2.
The inset shows the spectral region 275 nm to 325 nm. — No Fe(II),
— 1.00 mg/L Fe(II), — 2.00 mg/L Fe(II), — 5.00 mg/L Fe(II), —
9.99 mg/L Fe(II)
decrease of these two peaks indicates that the concentration of the “free” LGB reagent is lower, due
to the presence of the iron(II) ion. One possibility for this decrease is the complex ion formation of
LGB with iron. Based on the structure of LGB (Figure 4), it is not expected to be a strong
complexing agent. Citrate ion is known to form strong complexes with iron in both +2 and +3
oxidation states. The logarithms of the formation constant of the citrate complexes are 6.15 and 13.22
for iron(II) and iron(III), respectively. The logarithms of the formation constant of the hydrogen
citrate complexes are 10.2 and 14.45 for iron(II) and iron(III), respectively. The presence of a
complex between iron and LGB is even less likely in the presence of strong complexing agents, such
as the citrate ion. However, if the concentration of citrate ion is lower than the concentration of iron,
LGB may form a complex with iron. To rule out this possibility, the moles of citrate ion species and
iron in the reaction mixture were calculated based on the known dilutions. The total amount of iron
52
species is 4.4×10–3 mmol with 10 mg/L added iron(II). The total amount of citrate species is
0.72 mmol. Because citrate ion is a tridentate ligand, it forms primarily 1:1 complexes with iron.
Thus, the citrate ion species are present in high enough excess to complexate all iron species even at
the highest tested iron(II) concentration.
The other possibility for the decrease in the free LGB concentration is that it undergoes a redox
reaction with iron(II). Based on the current results this may be the most feasible explanation.
The results of the interference measurements are summarized in Table 9 and Figure 12.
Figure 12. The change of the determined ClO2 concentration with
iron(II) concentration.  no ClO2 added, Ž 0.25 mg/L ClO2 added, •
2.0 mg/L ClO2 added. The dashed line show 0.25 mg/L determined ClO2
concentration.
53
Table 9. Determined ClO2 concentrations in the presence of iron(II)
ion, in mg/L units.
True ClO2 (mg/L)
[Fe2+] (mg/L)
0
0.25
2
0
N/A
0.26 (±0.04)
1.71 (±0.05)
1
0.02 (±0.04)
0.09 (±0.04)
1.24 (±0.01)
2
0.09 (±0.02)
0.07 (±0.04)
0.78 (±0.02)
5
0.10 (±0.04)
0.15 (±0.03)
0.11 (±0.03)
9.99
0.15 (±0.06)
0.27 (±0.04)
0.13 (±0.03)
Average*
0.09
N/A
N/A
± represents the standard deviation of three samples
* in the case of blank solutions (no ClO2 added, second column) it is the
standard error from zero. This is calculated by averaging the absolute values of
the determined ClO2 concentrations.
When no ClO2 is added, in the presence of 1 and 2 mg/L iron(II), the measured absorbances are
not significantly different from the absorbance of the blank solution (<0.03 absorbance unit
difference). At higher iron concentrations, however, the absorbance difference exceeds the previously
defined ±0.03 absorbance unit difference. Thus, at these levels, iron is an interferent. Even though
at 5 and 10 mg/L iron concentration the determined ClO2 concentrations are in a region where the
accuracy of the measurements is low, an increase is observed in the determined concentration with
increasing iron concentration. This increase is possibly due to the reaction between iron(II) and LGB.
When 0.25 mg/L ClO2 is added to the iron(II) solution, the determined ClO2 concentrations at
1.00 and 2.00 mg/L iron(II) concentration are not significantly different from the blank solution. At
5.00 and 9.99 mg/L iron(II) concentration, the determined ClO2 concentration is above the detection
limit and increases with increasing iron(II) concentration. At 9.99 mg/L iron(II) concentration, the
determined ClO2 concentration is not significantly different from the true ClO2 concentration.
54
When 2.0 mg/L ClO2 is added to the Fe(II) solution, the determined ClO2 concentration
decreases with increasing iron(II) concentration up to 5 mg/L. At 5 mg/L iron concentration, a
possible minimum or plateau region is reached in the determined ClO2 concentration.
These observations can be understood by considering the following. There are two factors that
contribute to the changes in the determined ClO2 concentration in this case. One is the reaction
between LGB and iron(II). This process results in a decrease in the absorbance (positive error). The
error due to this reaction increases with increasing iron concentration. The magnitude of this error
can be seen from the blank experiments, which did not contain added ClO2.
The other process is the reaction between Fe(II) and ClO2. This reaction presents a demand for
ClO2. Iron(II) can be relatively easily oxidized by ClO2 that is a strong oxidizing agent. The
appropriate formal potentials19, 78 are given in Equations 24 and 25.
Fe3+ + e– ¾ Fe2+
0.771 V
(24)
ClO2 + e– ¾ ClO2–
1.16 V
(25)
ClO2– + 2 H2O + 4 e– ¾ Cl– + 4 OH–
0.76 V
(26)
Based on these potential values, it can be seen that ClO2 is able to oxidize iron(II) to iron(III)
and form chlorite ion. Chlorite ion can react further with iron(II) resulting in the following overall
reaction9.
ClO2 + 5 Fe2+ + 13 H2O Ž 5 Fe(OH)3 + Cl– + 11 H+
(27)
However, iron(III) can catalyze the decomposition of chlorite ion79, 80. This reaction decreases
ClO2 demand, because the decomposition products of chlorite ion can be ClO2, chlorate, and chloride
ions. In addition to these products, several short-lived, reactive intermediates are formed. The
55
intermediates may react with ClO2 or chlorite ion to form further chlorine containing species (e.g.,
hypochlorous acid). Another possibility is that the intermediates react with the LGB or the oxidation
products of the LGB. Thus, the iron(II)-ClO2 system is a complex reaction system, where the
composition of the final solution is greatly influenced by the initial parameters, such as the ratio of
ClO2 to Fe(II) and the chloride ion concentration. Some of the reactions which take place are
summarized below. The detailed reaction mechanism is given by Fábián and Gordon80.
Fe2+ + ClO2 Ž Fe3+ + ClO2–
(28)
Fe3+ + ClO2– ¾ FeClO22+
(29)
FeClO22+ ¾ Fe2+ + ClO2
(30)
Fe2+ + ClO2– Ž Fe3+ + Cl(II)
(31)
ClO2– + Cl(II) Ž HOCl + ClO2
(32)
The various chlorine containing species react further with each other. To describe the reaction
of these chlorine species, nine or ten additional equations are required. This well illustrates the
complexity of this system.
In the above reactions ClO2 is consumed. Thus, the determined ClO2 concentration is lower than
the added ClO2 concentration due to these reactions. The determined ClO2 concentration decreases
with increasing iron concentration. This decrease continues to the iron concentration at which all ClO2
is consumed. At higher concentrations, the determined ClO2 concentration would remain the same
if this were the only process which alters the determined ClO2 concentration.
The observed changes in the determined ClO2 concentrations are due both to interference process
and demand for ClO2. However, at high iron(II) concentration when all ClO2 is consumed, the change
56
in the ClO2 concentration is only dependent on the reaction between the LGB-iron(II). Thus, an
increase in the determined ClO2 concentration at high iron concentration would be observed.
However, the iron concentration at which all ClO2 is consumed, depends on the ClO2 concentration.
In the current measurements, the 0.25 mg/L ClO2 is probably used up even by 1 mg/L Fe(II). For
2 mg/L ClO2, this concentration is about 5 mg/L Fe(II) as indicated by the beginning of a plateau
region in the determined ClO2 concentration plot.
On the other hand, at low iron concentration, where not all ClO2 is used up, both processes
contribute to the changes in the determined ClO2 concentration. The observed changes in the
determined ClO2 concentration are dependent on the relative magnitude of these two sources of
errors.
Chiswell and O’Halloran39 in their original paper reported that oxidized forms of iron did not
present an interference. However, no experimental details are given. No reaction is expected between
ClO2 and iron(III) as iron(III) is not able to reduce or oxidize ClO2. Furthermore, iron(III) is not able
to oxidize LGB, due to the high redox potential of the dye.
Conclusions: The results suggest that iron(II) reacts with LGB. This reaction results in a
relatively low increase in the determined ClO2 concentration. Iron(II) reacts with ClO2, resulting in
a demand for ClO2. This reaction is not a direct interference with the measurement of ClO2 because
it takes place before sampling. Iron(III), the product of this reaction, is not an interference as it has
been shown by Chiswell and O’Halloran39.
57
2.6.6. Manganese(II) interference
The secondary drinking water standard77 for manganese (without any specification for its
oxidation state) is 0.05 mg/L. The interference of manganese(II) was studied in the 1–10 mg/L
concentration range. The possible interference mechanisms of manganese(II) are outlined below.
!
Reaction with LGB
"
Complex formation
"
Redox reaction
!
Complex formation with citrate and hydrogen citrate ions
!
Reaction with ClO2 (demand for ClO2)
The stoichiometric oxidation product of manganese(II) is manganese dioxide. Manganese dioxide
can not be reduced by chlorite ion. Rather, chlorite ion oxidizes manganese(II). In addition,
manganese dioxide can not be oxidized to permanganate ion due to the high redox potential of the
permanganate ion.
Even though manganese(II) forms strong complexes with citrate and hydrogen citrate ions, no
spectral changes were observed in the UV region of the measured spectra that would be indicative
of these species. The logarithms of the formation constants of the manganese(II) complexes are 5.5
and 9.6 with citrate and hydrogen citrate ions, respectively. Thus, the presence of a complex between
manganese(II) and LGB can be excluded on the same basis as in the case of iron(II).
The results of the measurements are shown in Table 10 and Figure 13. The measured absorbance
was within ±0.03 absorbance units from the absorbance of the double blank solution (reagent water)
in every measurement when manganese(II) was added. Based on these results, two conclusions can
be drawn. First, manganese(II) does not react with LGB. If there were a reaction between LGB
58
Table 10. Determined ClO2 concentrations in the presence of
Mn(II), in mg/L units
Added ClO2 (mg/L)
[Mn2+]
(mg/L)
0
0.25
2
0
N/A
0.16 (±0.03)
1.98 (±0.11)
1
0.12 (±0.01)
0.13 (±0.02)
0.17 (±0.04)
2
0.15 (±0.01)
0.09 (±0.04)
0.13 (±0.02)
5
0.10 (±0.03)
0.15 (±0.04)
0.19 (±0.02)
10
0.17 (±0.06)
0.04 (±0.06)
0.16 (±0.05)
average*
0.13
0.1
0.16
corrected**
N/A
-0.03
0.03
± represents the standard deviation of three samples
* in the case of blank solutions (no ClO2 added, second column) it is the
standard error from zero. This is calculated by averaging the absolute values of
the determined ClO2 concentrations.
**calculated by subtracting the average determined ClO2 concentration for blank
solution from the average determined ClO2 concentration
Figure 13. The change of the determined ClO2 concentration
with Mn(II) concentration.  No ClO2 added, Ž 0.25 mg/L
ClO2 added, • 2.0 mg/L ClO2 added
59
and manganese(II), the difference between the absorbances of the blank solution and the
manganese(II) containing solutions would be expected to be higher than 0.03 absorbance units.
Furthermore, if manganese(II) reacted with LGB, a trend would be expected in the determined ClO2
concentration versus the manganese(II) concentration.
The second conclusion is that manganese(II) reacts with ClO2, resulting in a demand for ClO2.
This is illustrated by the low determined ClO2 concentrations when 2.0 mg/L ClO2 is added to the
manganese(II) solutions. When 0.25 mg/L ClO2 is added to the manganese(II) solutions, the
determined ClO2 concentrations, in the 1 to 5 mg/L manganese(II) concentration range, are
statistically not different from the determined ClO2 concentration in the absence of manganese(II).
However, the results of the other two series of measurements indicate that these measured ClO2
concentrations are not truly due to the added ClO2.
Even though, the determined ClO2 concentrations are considerably higher than zero in the
absence of added ClO2, the measured absorbance is not significantly different from the absorbance
of the blank solution. This is indicated by the absorbance difference which is below ±0.03 absorbance
units (three-times the standard deviation of the blank solutions).
Similar to iron(II), manganese(II) can be oxidized by ClO2 according to the following equation9.
2 ClO2 + 5 Mn2+ +6 H2O Ž 5 MnO2(s) + 12 H+ + 2 Cl–
(33)
According to this equation, manganese(II) can effectively remove ClO2. In this case, the
stoichiometric ratio of ClO2 to manganese (2:5) is higher than for iron (1:5). This means that less
manganese(II) is needed to remove 2 mg/L ClO2 than iron(II). This is illustrated by the low
determined ClO2 concentrations even when only 1 mg/L manganese(II) is present.
60
No interference process is present in the case of manganese(II). The only process that takes place
in the presence of manganese(II) is the demand for ClO2 by the manganese(II)–ClO2 reaction.
Because of the stoichiometry of the manganese(II) and ClO2 reaction, the determined ClO2
concentration is low and very close to the detection limit in the presence of even 1 mg/L
manganese(II).
Conclusions: Manganese(II) has a significant demand for ClO2, due to the reaction between ClO2
and Mn(II) that results in practically complete removal of ClO2. This reaction takes place before
sampling in real measurements. Thus, it does not present a direct interference with the measurement
of ClO2.
2.6.7. Manganese(VII) interference
To test the interference of manganese(VII) (permanganate ion) on the measurement of ClO2 with
LGB method, potassium permanganate solutions were used in the 1–10 mg/L concentration range.
The potassium permanganate stock solution was standardized by using sodium oxalate67. The
expected interference mechanisms can be summarized as follows.
!
Oxidation of LGB
!
Oxidation of ClO2 (demand for ClO2)
!
Absorbance increase at 633 nm
The absorbance of the permanganate solutions was measured. The absorbance of the 10 mg/L
permanganate ion solution at 633 nm was 7×10–3 absorbance units. This is below the standard
deviation of the absorbance measurement of the blank solutions. Thus, permanganate ion solutions
do not change the absorbance significantly at 633 nm.
61
The results are shown in Table 11 and Figure 14. When permanganate ion solution is added to
the LGB solution, significant absorbance decrease is observed both in the absence and presence of
Table 11. Determined ClO2 concentrations in the presence of
permanganate ion, in mg/L units.
Added ClO 2 (mg/L)
[MnO 4 – ]
(mg/L)
0
0.25
2
0
N/A
0.16 (±0.03)
1.98 (±0.11)
1
0.53 (±0.03)
0.51 (±0.05)
1.61 (±0.02)
2
1.02 (±0.06)
0.97 (±0.05)
1.44 (±0.03)
5
1.90 (±0.03)
2.00 (±0.04)
2.04 (±0.02)
10
2.82 (±0.03)
2.99 (±0.01)
2.84 (±0.02)
± represents the standard deviation of three samples
Figure 14. The change of the determined ClO2 concentration
with permanganate ion concentration. The lines show the least
square fit of the data.  No ClO2 added, Ž 0.25 mg/L ClO2
added, • 2.0 mg/L ClO2 added. The equation of the lines: —
[ClO2]det. = 0.245×[MnO4–] + 0.464, R2 = 0.970, — [ClO2]det.
= 0.270×[MnO4–] + 0.402, R2 = 0.972
62
ClO2. When no ClO2 or 0.25 mg/L ClO2 is added, the observed absorbance decrease is linearly
dependent on the permanganate ion concentration as the least squares fit shows in Figure 14. When
2.0 mg/L ClO2 is added, after an initial decrease the determined ClO2 concentration increases with
increasing permanganate ion concentration.
The determined ClO2 concentrations in the presence of 10 mg/L permanganate ion are outside
of the calibration range of the ClO2 concentrations. These values were determined by extrapolation
and for this reason their accuracy is lower than the other determined concentrations.
The change of the determined ClO2 concentration with permanganate ion concentration was
compared when no ClO2 or 0.25 mg/L ClO2 was added. The two lines were compared by using
multiple model linear regression. The tentative model contained three independent variables: the
permanganate ion concentration, an indicator variable for the presence of 0.25 mg/L ClO2 and the
product of the previous two. The indicator variable is 0 when no ClO2 is present and 1 when
0.25 mg/L ClO2 is present. This model assumes that the variances of the error terms are the same but
permits the two regression lines to have different slopes and intercepts. A fit of this regression model
to the data showed that at the 95% confidence level neither the indicator variable nor the product of
the indicator variable and the permanganate ion concentration are significant. Thus it is concluded
that neither the intercepts nor the slopes of the two lines are statistically different.
Permanganate ion is a strong oxidizing agent with a redox potential that is significantly higher
than the redox potential of LGB. This means that permanganate ion can oxidize LGB, resulting in an
absorbance decrease. This absorbance decrease is linearly dependent on the concentration of
permanganate ion as the results show (see Figure 14).
63
In addition, permanganate ion is stronger oxidizing agent than ClO2. Thus, permanganate ion
oxidizes ClO2, forming chlorate ion. When 0.25 mg/L ClO2 is added to the permanganate ion solution,
all ClO2 is oxidized as indicated by the statistically indistinguishable linear fits in the absence of ClO2
and in the presence of 0.25 mg/L ClO2.
The non-linear change of the determined ClO2 concentration at 2.0 mg/L added ClO2
concentration is due to the fact that not all ClO2 is oxidized at low permanganate ion concentrations.
This curve can help to determine the stoichiometry of the ClO2–permanganate ion reaction. The
possible reactions are summarized below.
5 ClO2 + MnO4– + H2O Ž 5 ClO3– + Mn2+ + 2 H+
(34)
3 ClO2 + MnO4– + H2O Ž 3 ClO3– + MnO2(s) + 2 H+
(35)
ClO2 + MnO4– + H2O Ž ClO3– + MnO42– + 2 H+
(36)
To determine the stoichiometry, three additional measurements were performed at 3.0, 3.5, and
4.0 mg/L permanganate ion concentrations. Figure 15 shows the determined ClO2 concentrations as
a function of the permanganate ion to ClO2 molar ratio.
Figure 15. The change of the determined ClO2 concentration
with the permanganate ion to ClO2 ratio.
64
The minimum of this curve appears to be around 0.8 permanganate ion to ClO2. This is not
consistent with any of the above Equations and suggests that a combination of these reactions takes
place. The first step of the ClO2 permanganate ion reaction is Equation 36. If ClO2 is in excess,
Equation 35 also takes place to some extent. Equation 36 is a fast reaction, because it is a one
electron transfer reaction. This would result in a 1:1 permanganate ion to ClO2 ratio. Equation 35 is
a much slower reaction, because the transfer of three electrons is accompanied by significant
structural changes. By itself, this reaction would result in 1:3 permanganate ion to ClO2 ratio. The
observed 0.8 ratio shows that a combination of these two reactions takes place.
Conclusion: Permanganate ion is able to oxidize LGB, resulting in severe interference. However
at these permanganate ion concentrations the water is colored due to the presence of permanganate
ion. This is not acceptable in potable water, thus permanganate ion would be removed in the water
treatment plant. In addition, permanganate ion has a demand for ClO2. This demand would not
interfere with the measurement of ClO2 as the probable product of the ClO2–permanganate ion
reaction is manganese dioxide, which has been shown not to interfere39.
2.6.8. Manganese(II)–Manganese(VII) interference
It is known that in certain cases, the reduction of permanganate ion is catalyzed by
manganese(II). One example is the standardization of potassium permanganate solutions with sodium
oxalate. To increase the rate of this reaction, the titrated solution is heated (about 70-80 °C), and a
small amount of manganese(II) salt can be added to increase the rate of the reduction of
permanganate ion.
65
Therefore, it is necessary to study the interference of mixtures of manganese(II) and
permanganate ion. The manganese(II) concentration was varied between 0.1 mg/L and 10 mg/L. The
permanganate ion concentration was varied between 1 mg/L and 10 mg/L.
Upon mixing manganese(II) solutions with permanganate ion solutions, a color change was
observed to red-orange. After a few minutes, the color became lighter and a dark brown precipitate
was observed. This precipitate is probably manganese dioxide, formed according to Equation 37.
3 Mn2+ + 2 MnO4– + 2 H2O Ž 5 MnO2 + 4 H+
(37)
The possible interference processes are summarized below.
!
!
Permanganate ion reactions
"
Reaction with LGB
"
Reaction with ClO2 (demand for ClO2)
Manganese(II) reactions
"
Reaction with ClO2 (demand for ClO2)
The various other interference processes have been ruled out in the description of the individual
interferents. These three interferent processes are not present simultaneously, due to Equation 37.
Thus, either the permanganate or the manganese(II) interference processes are observed, depending
on the ratio of permanganate to manganese(II). Table 12 and Figures 16-18 show the results of these
experiments.
66
Table 12. Determined ClO2 concentrations in the presence of permanganate ion
and manganese(II), in mg/L units.
Added ClO2 (mg/L)
[Mn2+]
(mg/L)
[MnO4–]
(mg/L)
0
0.25
2
0
0.00 (day 1)*
N/A
0.28 (±0.08)
1.93 (±0.02)
0
0.00 (day 2)*
N/A
0.25 (±0.03)
1.95 (±0.00)
0.1
1
0.43 (±0.02)
0.65 (±0.03)
1.51 (±0.04)
0.1
2
0.87 (±0.04)
1.16 (±0.08)
1.32 (±0.05)
0.1
5
1.85 (±0.01)
2.48 (±0.03)
1.89 (±0.02)
0.1
10
2.84 (±0.02)
3.88 (±0.03)
2.74 (±0.04)
1
1
0.16 (±0.00)
0.04 (±0.05)
0.93 (±0.05)
1
2
0.44 (±0.03)
0.57 (±0.10)
1.03 (±0.01)
1
5
1.62 (±0.04)
2.12 (±0.05)
1.72 (±0.02)
1
10
2.85 (±0.04)
3.83 (±0.04)
2.65 (±0.02)
5
1
0.10 (±0.04)
0.14 (±0.02)
0.10 (±0.02)
5
2
0.13 (±0.05)
0.15 (±0.04)
0.09 (±0.02)
5
5
0.15 (±0.02)
0.15 (±0.06)
0.18 (±0.04)
5
10
1.44 (±0.05)
1.56 (±0.05)
1.97 (±0.05)
10
1
0.17 (±0.02)
0.11 (±0.05)
0.13 (±0.04)
10
2
0.14 (±0.05)
0.07 (±0.02)
0.12 (±0.05)
10
5
0.12 (±0.08)
0.17 (±0.06)
0.09 (±0.02)
10
10
0.03 (±0.06)
0.15 (±0.01)
0.11 (±0.03)
± represents the standard deviation of three samples
* Each set of measurements was completed on two days.
67
Figure 16. The change of the determined ClO2 concentration
with permanganate ion concentration in the presence of
Mn(II). No ClO2 was added.  0.1 mg/L Mn(II), Ž 1.0 mg/L
Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
Figure 17. The change of the determined ClO2 concentration
with the permanganate ion concentration in the presence of
Mn(II). 0.25 mg/L ClO2 was added.  0.1 mg/L Mn(II), Ž
1.0 mg/L Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
68
Figure 18. The change of the determined ClO2 concentration
with permanganate ion concentration in the presence of
Mn(II). 2.0 mg/L ClO2 added.  0.1 mg/L Mn(II), Ž 1.0 mg/L
Mn(II), • 5.0 mg/L Mn(II), – 10.0 mg/L Mn(II)
The figures show that upon the addition of increasing amounts of manganese(II), the determined
ClO2 concentration becomes lower at the same permanganate ion and added ClO2 concentration. The
results indicate that the measurement of ClO2 is affected by mixtures of manganese(II) and
permanganate ion solutions in a complex manner. The primary parameter, which determines the
magnitude of this effect, is the ratio of manganese(II) to permanganate ion. Figure 19 shows the
change of the determined ClO2 concentration with the ratio of manganese(II) to permanganate ion.
If more than stoichiometric amount of permanganate ion is present, the effect is similar to the
permanganate interference. In this case, at the same Mn2+/MnO4– ratio, the determined ClO2
concentration is also dependent on the permanganate ion concentration. If manganese(II) is present
in higher than stoichiometric concentration, the observed interference is similar to the manganese(II)
case. In this case, independent of the added ClO2 concentration, the determined ClO2 concentration
is at the detection limit.
69
Figure 19. The change of the determined ClO2 concentration
with the manganese(II) to permanganate ion molar ratio. 
1.0 mg/L MnO4–, Ž 2.0 mg/L MnO4–, • 5.0 mg/L MnO4–, –
10.0 mg/L MnO4–
Conclusion: Mixtures of permanganate ion and manganese(II) interfere with the measurement
of ClO2 in a complex manner. When permanganate ion is in excess, the determined ClO2
concentrations are high and due to the oxidation of LGB by permanganate. In the case of excess
manganese(II), the determined ClO2 concentrations are low due to the reaction between
manganese(II) and ClO2. However, no indication was found that the presence of manganese(II) acts
as a catalyst for the permanganate ion reaction. The only effect of Mn(II) on the permanganate ion
interference is by means of Equation 37, which reduces the concentration of permanganate ion. This
lower permanganate ion results in lower determined ClO2 concentration.
70
2.6.9. Free Available Chlorine (FAC) interference
Chlorine can be used either before or after the addition of ClO2 to the water to improve the
disinfection. The MRDL for dissolved chlorine6, 7 is 4.0 mg/L. The tested concentration range was
from 1.0 mg/L to 8.0 mg/L. The possible interference processes for FAC are summarized below.
!
Reaction with LGB
!
Reaction with ClO2 (demand for ClO2)
!
Reaction of chloroaminoacetic acid with LGB (chloroaminoacetic acid is the product of
the reaction between glycine and FAC.)
However, glycine is expected to mask FAC by forming chloroaminoacetic acid. Thus, no reaction
is expected between LGB and FAC. The amount of glycine present is 0.17 mmol. The amount of
FAC at 8 mg/L concentration is 2.77×10–3 mmol. This means that a large enough excess of glycine
is present to mask FAC interference even at the highest tested concentration. The results are shown
in Figure 20 and Table 13.
Figure 20. The change of the determined ClO2 concentration
with FAC concentration.  no ClO2 added, Ž 0.25 mg/L ClO2
added, • 2.0 mg/L ClO2 added
71
Table 13. Determined ClO2 concentrations in the presence of FAC,
in mg/L units.
Added ClO2 (mg/L)
[FAC]
(mg/L)
0
0.25
2
0
N/A
0.27 (±0.02)
2.04 (±0.04)
1
-0.02 (±0.01)
0.25 (±0.04)
2.01 (±0.04)
4
0.00 (±0.00)
0.21 (±0.04)
2.04 (±0.01)
6
0.00 (±0.02)
0.29 (±0.02)
1.98 (±0.02)
8
0.01 (±0.01)
0.27 (±0.02)
1.79 (±0.04)
Average*
0
0.25
1.96
Corrected**
N/A
0.25
1.96
± represents the standard deviation of three samples
* in the case of blank solutions (no ClO2 added, second column) it is the
standard error from zero. This is calculated by averaging the absolute values
of the determined ClO2 concentrations.
**calculated by subtracting the average determined ClO2 concentration for
blank solution from the average determined ClO2 concentration
When no ClO2 is added to the FAC solution, the determined absorbances are not significantly
different from the determined absorbance of the blank solution. When 0.25 mg/L ClO2 is added, the
determined ClO2 concentrations are not significantly different from the ClO2 concentration which was
determined in the absence of FAC. These results show that neither FAC, nor chloroaminoacetic acid
reacts with LGB.
When 2.0 mg/L ClO2 is added to the FAC solution in the 1.0 mg/L to 6.0 mg/L concentration
range, the determined ClO2 concentrations are not significantly different from the ClO2 concentration
in the absence of FAC. However, when 8.0 mg/L FAC is present, the determined ClO2 concentration
is lower than the added concentration. At these concentration levels, the probable reason for the
decrease in the ClO2 concentration is the reaction between dissolved chlorine and ClO2. This
72
reaction17, 18 presents a demand for ClO2. The reaction is moderately fast and is more important at
high concentrations than at low concentrations.
Conclusions: The results suggest that glycine masks FAC efficiently. Thus, no interference is
observed from FAC at any of the tested FAC concentrations. However, a small decrease was
observed in the determined ClO2 concentration, due to the demand for ClO2 by FAC.
2.6.10. Monochloramine (NH 2Cl) interference
When chlorine is added to water that contains ammonia, various chloramines are formed,
depending on the pH and chlorine to ammonia ratio. Thus chloramines may be present in ClO2 treated
water that was subject to the combination of chlorine and ClO2. Furthermore, chloramines may be
added directly to ClO2 treated water. Chloramines are added to treated water to maintain residual
disinfectant concentrations throughout the distribution system2. Chloramines are generally formed at
the point of application by mixing chlorinated water with ammonia containing raw water. This
mixture is added to the treated water. The MRDL for chloramines6, 7 is 4 mg/L (as Cl2). The tested
concentration range was from 1 to 8 mg/L.
In water treatment practice, the concentration of the various disinfectant species is given in terms
of chlorine equivalent concentration or “as Cl2.” For example in the case of monochloramine, 1 mg/L
(as Cl2) monochloramine can be interpreted as follows. This is the concentration of monochloramine
that is able to transfer the same moles or number of electrons as 1 mg/L Cl2. Both Cl2 and
monochloramine undergo a 2 electron reduction. Thus, 1 mg/L (as Cl2) monochloramine is:
(38)
73
The results in Table 14 and Figure 21 show that even in the absence of ClO2, the determined
ClO2 is significantly different from zero and comparable to the concentration which is determined
when 0.25 mg/L ClO2 is added. This means that the method would falsely measure low
concentrations of ClO2 in the presence of monochloramine.
When 0.25 mg/L ClO2 is added to the chloramine solution, the determined ClO2 concentrations
are not significantly different from the ClO2 concentration determined in the absence of chloramine.
When 2.0 mg/L ClO2 is added, the determined ClO2 concentrations are significantly different at 1 to
4 mg/L chloramine concentrations, from the determined ClO2 concentration in the absence of
chloramine. However, these concentrations are still within the acceptable range (100 ±30%) defined
by US EPA. The determined ClO2 concentrations at 6 and 8 mg/L chloramine concentration are not
significantly different from the ClO2 concentration in the absence of chloramine.
Conclusions: Monochloramine interferes with the measurement of ClO2 because it results in a
false, low ClO2 concentration even in the absence of ClO2. Thus, it would be necessary to confirm
the presence of ClO2 by using some other method. On the other hand, the determined ClO2
concentrations are not influenced by the presence of the chloramine. In the current form the LGB
method can not be accurately used to measure ClO2 if monochloramine is present.
2.6.11. Conclusions on the interference results
The results of the interference study show that iron(II), manganese(II), permanganate ion, and
the mixture of manganese(II) and permanganate ion present a demand for ClO2. In addition to these
species, FAC has a demand for ClO2 at 8 mg/L FAC concentration. The demand of these interferents
can be easily misinterpreted as interference with the measurement of ClO2. The reactions between
these interferents take place during water treatment before sampling. Thus, the decrease in
74
Table 14. Determined ClO2 concentrations in the presence of
monochloramine, in mg/L units.
Added ClO2 (mg/L)
[NH2Cl]
(mg/L)
0
0.25
2
0
N/A
0.27 (±0.02)
2.04 (±0.04)
1
0.22 (±0.04)
0.25 (±0.05)
1.88 (±0.05)
4
0.21 (±0.04)
0.25 (±0.04)
1.90 (±0.04)
6
0.15 (±0.04)
0.29 (±0.01)
1.97 (±0.07)
8
0.19 (±0.02)
0.28 (±0.02)
2.03 (±0.08)
Average*
0.19
0.27
1.94
Corrected**
N/A
0.08
1.76
± represents the standard deviation of three samples
* in the case of blank solutions (no ClO2 added, second column) it is the
standard error from zero. This is calculated by averaging the absolute values
of the determined ClO2 concentrations.
**calculated by subtracting the average determined ClO2 concentration for
blank solution from the average determined ClO2 concentration
Figure 21. The change of the determined ClO2 concentration
with monochloramine concentration.  no ClO2 added, Ž
0.25 mg/L ClO2 added, • 2.0 mg/L ClO2 added
75
the ClO2 concentration due to these species is not an analytical problem, rather water treatment
technological difficulty.
Monochloramine presents a different problem. It does not alter the added ClO2 concentration,
but it results in a small measured ClO2 concentration in the absence of ClO2.
Table 15 gives a summary of the results in the absence of added ClO2. Iron(II) and permanganate
ion react with LGB, resulting in non-zero determined ClO2 concentrations even in the absence of the
analyte (see Table 15). This apparent ClO2 concentration linearly changes with the interferent in both
cases. The reaction of LGB with iron(II) presents only a relatively small interference, at 10 mg/L
iron(II) concentration the apparent ClO2 concentration is about 0.1 mg/L. On the other hand,
interference from permanganate ion is severe, 1 mg/L permanganate ion results in the measurement
of about 0.25 mg/L apparent ClO2 concentration.
Chlorate ion and dissolved chlorine do not interfere with the measurement of ClO2. These
compounds are generally encountered as an interference during in the measurement of ClO2 in potable
water. In many of the currently existing analytical methods, these compounds interfere and their
interference can be masked only by complicated procedures. For example, the DPD method9, 43 can
differentiate between chlorine, ClO2, and monochloramine only by means of a stepwise procedure.
Thus, the currently proposed procedure for the LGB method is selective for ClO2 in the presence
of chlorine and chlorate ion. However, iron and permanganate interfere with the measurements. It
has been shown earlier35, 36, 49, 81 that the selectivity of ClO2 analytical methods can be improved by gas
diffusion flow injection analysis (GD-FIA). The GD-FIA method is expected to improve the
selectivity of the current LGB method.
76
Table 15. Summary of the measured ClO2 concentrations in the presence of
various species, in mg/L units. No ClO2 was added to the solutions.
ClO3–
Fe2+
Mn2+
MnO4–
FAC
NH2Cl
1 mg/L
-0.01
(±0.04)
0.02
(±0.04)
0.12
(±0.01)
0.53
(±0.03)
-0.02
(±0.01)
0.22
(±0.04)
2 mg/L
-0.09
(±0.14)
0.09
(±0.02)
0.15
(±0.01)
1.02
(±0.06)
N/A
N/A
4 mg/L
N/A
N/A
N/A
N/A
0.00
(±0.00)
0.21
(±0.04)
5 mg/L
0.04
(±0.01)
0.10
(±0.04)
0.10
(±0.03)
1.90
(±0.03)
N/A
N/A
6 mg/L
N/A
N/A
N/A
N/A
0.00
(±0.02)
0.15
(±0.04)
8 mg/L
N/A
N/A
N/A
N/A
0.01
(±0.01)
0.19
(±0.02)
10 mg/L
0.03
(±0.04)
0.15
(±0.06)
0.17
(±0.06)
2.82
(±0.03)
N/A
N/A
± represents the standard deviation of three samples
2.7. Conclusions
In order to promote the use of ClO2, the US EPA has developed a new LGB method (proposed
EPA Method 327.0). The purpose of this method is to provide a simple, reliable analytical method
for compliance monitoring.
The proposed EPA Method 327.0 measures both ClO2 and chlorite ion in treated water in the
0.25 to 2.0 mg/L range. This range can be extended by adjusting sample and reagent volumes in order
to measure lower or higher concentrations. The method shows good linearity in the 0.25 to 2.0 mg/L
ClO2 or chlorite ion concentration range. The LGB method works well in both reagent and tap water.
The accuracy of the measurements of a single analyte is good.
77
The proposed Method 327.0 can accurately determine ClO2 concentrations in the presence of
chlorine, chloramine, and chlorate ion. Iron(II), manganese(II), and permanganate ion interfere with
the method. The interference from these species can be eliminated either by masking them or by using
gas diffusion flow injection analysis.
2.8. Future directions
This section summarizes the problems (i.e., shortcomings) encountered during the second
laboratory and interference studies. The purpose of summarizing these problems is to make
suggestions for possible solutions. In addition, potential improvements are suggested which would
make the current method preferable over other analytical methods.
2.8.1. Chlorine dioxide standards
A general problem with ClO2 analytical methods is that due to the volatile, reactive nature of
ClO2, no “factory made” standard solutions can be prepared. The calibration of these methods would
be easier if standard ClO2 solutions were available with accurately known composition, without the
need to standardize a stock ClO2 solution and prepare the necessary dilutions.
Preliminary experiments have shown that it may be possible to create reliable, stable ClO2
standards by mixing chlorite ion solutions with an appropriate photoacid and illuminating the mixture
with a suitable light source. Photoacids are special organic compounds that become more acidic upon
photo excitation82, 83. Considering acid AH, the dissociation reactions in the ground and excited states
are:
AH ¾ A– + H+
78
(39 a)
*AH ¾ *A– + H+
(39 b)
The asterisk denotes the excited molecules. The pKa and pKa* are the acid dissociation constants
of the ground state acid molecule and the excited state acid molecule, respectively. The relationship
between pKa and the free energy is
(40 a)
(40 b)
The increase in the acidity of the organic acid in the excited state is
(41)
where <1
<2
the excitation frequency of the acidic form
the excitation frequency of the basic form
This equation is called Förster equation83, 84. Based on this equation, the excited state acid
molecule is stronger acid if the excitation frequency of the basic form is lower than the excitation
frequency of the acidic form (bathochromic shift of the emission or absorption spectrum)83.
There is a significant number of organic dyes that undergo significant color change upon
deprotonation. These dyes are possible candidates for being photoacids. However, it is important that
these compounds easily undergo the proton transfer reaction83.
Some considerations are summarized here that are important in creating ClO2 standards by using
photoacid/chlorite ion mixtures. Chlorine dioxide is known to decompose upon exposure to light. For
this reason, it is important that the required illumination time or light intensity is high enough to
generate ClO2, but at the same time low enough to avoid the photodecomposition of ClO2. It is
important that the photoacid does not react with chlorite ion directly and that only the acid-base
79
reaction takes place. This would ensure the long shelf-life of the ClO2 standard. The photoacid in the
excited state needs to have lower pKa value than chlorous acid (pKa = 1.72).
Upon illuminating the mixture of the photoacid and chlorite ion by a suitable light source, the
photoacid is converted to the excited, more acidic form. This strong acid would protonate chlorite
ion to form chlorous acid. As it has been shown (Equations 9 and 10) chlorous acid undergoes rapid
decomposition, forming ClO2.
4 HClO2 Ž 2 ClO2 + ClO3– + Cl– + 2 H+ + H2O
(9)
5 HClO2 Ž 4 ClO2 + Cl– + H+ + 2 H2O
(10)
The amount of ClO2 that forms in this sequence can be adjusted by changing the illumination time,
chlorite ion, and photoacid concentrations. In preliminary experiments, the photoacid 8-hydroxy1,3,6-trisulfonate-pyrene was used. The proposed method works, but the currently used photoacid
reacts directly with chlorite ion. This results in a short life-time of the mixture of chlorite ion and the
photoacid. This method appears to be able to generate accurately controlled concentrations of ClO2.
The advantage of this method is that ClO2 can be generated in a reproducible manner without
additional chemical expertise.
2.8.2. Using gas diffusion flow injection analysis with proposed EPA Method 327.0
It has been demonstrated previously9, 35-37, 49, 81 that by using gas diffusion flow injection analysis,
it is possible to minimize the interferences in the measurement of gaseous analytes in aqueous
solutions. The crucial part of GD-FIA is the gas diffusion cell. The two compartments of this cell are
separated by a porous membrane that is permeable only for gases. The sample is injected into a donor
stream. This donor stream passes through one compartment of the gas diffusion cell. The receiver
80
stream flows through the other compartment. The parameters (pH, reagent, etc.) of the receiver
stream can be adjusted to increase the gas transfer from the donor stream. After passing through the
gas diffusion cell, the receiver stream can be mixed with the necessary reagents to determine the
analyte concentration by measuring the absorbance change of the reagents.
The GD-FIA method can be used to improve the selectivity of the current method. It is a good
way to automate the measurement of ClO2 and chlorite ion. Furthermore, GD-FIA can be used to
eliminate the accuracy problems arising from the differential nature of the method. This could be
accomplished in the following way.
The sample would be injected into a pH 6.0 buffered donor stream. This stream would pass
through the gas diffusion cell where ClO2 would diffuse through the membrane into the receiver
stream. The receiver stream is a buffered LGB solution. After a short reaction time, ClO2 would react
with LGB and the absorbance would be measured at 633 nm. Chlorite ion, the other analyte, would
remain in the donor stream which would be mixed with the HRP solution. The chlorite ion would be
allowed to react with HRP for 20-30 minutes. The donor stream would pass through another gas
diffusion cell. In this cell, ClO2 formed from chlorite ion would diffuse through the membrane into
an LGB containing receiver solution (different from the previous receiver solution). The absorbance
decrease of this receiver solution would give the concentration of chlorite ion. Thus, the use of the
GD-FIA method potentially can make the LGB method better by improving reproducibility and
minimizing problems from interfering species.
81
3. The Cl2O 4– Complex
The application of ClO2 as a disinfectant in water treatment is greatly helped by simple analytical
methods, which can give accurate, real-time results, and are easy to perform. The spectrophotometric
measurement meets these criteria. This method gives instantaneous and accurate results, and can be
simply used for continuous measurement of ClO2. In fact, many currently existing ClO2 process
analyzers use photometric measurement for the determination of ClO2 concentrations.
The maximum ClO2 concentration that can be determined at the wavelength of the maximum
absorption (360 nm) is about 1.60x10–3 M (110 mg/L) by using a 1 cm cell. The concentrations of
ClO2 solutions in generator effluents (where on-line analyzers are necessary) are generally on the
order of few g/L. Thus, this high concentration needs to be measured at longer wavelengths where
ClO2 has a lower molar absorptivity or by using shorter pathlength cells. However, the trend among
the manufacturers of the analyzers is to use longer wavelengths instead of shorter pathlengths. The
wavelength used can be as high as 450 nm or even as high as 500 nm.
However, significant interference is caused indirectly by chlorite ion, due to a complex formation
reaction. The composition of the complex formed is Cl2O4–. The absorption spectrum of this complex
significantly overlaps with the spectrum of ClO2 and is shifted to longer wavelengths. This means that
above about 400 nm, the Cl2O4– complex has higher molar absorptivity than ClO2, which results in
increased absorbance in the presence of the complex. The increased absorbance results in false high
ClO2 concentrations. This false ClO2 concentration increase may lead to lower ClO2 doses than
required for the disinfection of potable water and possibly result in inadequate disinfection.
82
The objective of this research was to determine accurately the formation constant and molar
absorptivity of the Cl2O4– complex by using numerical methods. Based on these values, practical
methods and calculations can be devised to eliminate the interference of chlorite ion on the
spectrophotometric measurement of ClO2.
3.1. Theoretical
The details of the spectrophotometric measurement are given in Chapter 2. Here only the
relevant properties of this method are briefly summarized. The spectrophotometric measurement
presents a simple, fast, and reliable method for measuring ClO2 concentrations in a wide range. This
method can be used for measuring a variety of samples, for example measuring the ClO2
concentration in the field by using a pocket colorimeter or for measuring ClO2 in the generator
effluent. The maximum concentration of ClO2, which can be determined by using this method,
depends on the path length of the cell and on the wavelength at which the measurements are taken.
High ClO2 concentrations are generally measured at higher wavelengths than 360 nm. This
wavelength can be as high as 450 nm or even 500 nm, because in this region the molar absorptivity
of ClO2 is low.
3.1.1. The history of the Cl2O 4– complex
Bray noted first that if ClO2 is bubbled through a chlorite ion solution, the color of the solution
turns much darker than the color of the concentrated ClO2 solution, reaching an almost mahogany
color85. His PhD student, Barnett, observed that if ClO2 is bubbled through the solution for extended
periods of time, the brown color slowly changes to the dark yellow color of the concentrated ClO2
83
solution86. If ClO2 is removed, the remaining solution had no oxidizing capability, indicating that
chlorite ion has been decomposed. At that time he was unable to determine the composition of the
species formed due to the lack of sufficiently sophisticated and rapid analytical methods.
In another study, Gordon and Emenegger87 determined the composition of the complex that
formed in the system by using Job’s method88 and described its formation with the following equation:
ClO2 + ClO2– ¾ Cl2O4–
Keq = 1.6 M–1
(42 a)
(42 b)
Crawford, in his PhD thesis89, redetermined the formation constant (Keq) by minimizing the error
between the calculated and measured absorbances. In addition, he also determined the molar
absorptivity of the Cl2O4– complex. His results suggest that the complex absorbs in a similar region
as ClO2, but the spectrum of the Cl2O4– complex is slightly shifted toward longer wavelengths.
The reason for redetermining the formation constant and molar absorptivity is to determine more
accurate values and to give more chemical detail about the complex. Having more accurate values
than previously published would allow water treatment utilities and researchers to make appropriate
adjustments in the spectrophotometric measurement of ClO2 in order to maximize the benefits of
using ClO2 as a disinfectant.
The accuracy of the determined molar absorptivity and formation constant are greatly aided by
the improvements in the laboratory instrumentations and computers. To determine both the molar
absorptivity and the formation constant, numerical methods are needed which require a large number
of data points to improve their accuracy. By using diode array spectrophotometers, the collection of
84
this large amount of data can be easily accomplished. Furthermore, the manipulation of this quantity
of data is greatly aided by today’s high speed computers.
Even though the Cl2O4– complex has been known for almost a century, its existence is not widely
accepted or known. However, this situation is slowly changing, which is well illustrated in the
Handbook of Chlorination and Alternative Oxidants by White. In the third edition of this book90, the
existence of the complex is not referenced, but in the fourth edition2 the possible interference from
the Cl2O4– complex is discussed.
Despite the fact that the complex is becoming more widely known, its effects are not yet fully
appreciated or corrected for. One of the major reasons for the slow acceptance is probably the
relatively inaccurate value of the formation constant and molar absorptivity. The reasons for the
inaccuracy and large error in these values can be summarized as follows:
!
The low value of the formation constant means that even in concentrated solutions the relative
concentration of the complex is low.
!
The spectra of the Cl2O4– complex and ClO2 overlap.
!
Mathematically, the molar absorptivity of the complex and the value of Keq are not
independent variables.
3.2. Numerical methods
3.2.1. Matrix Rank Analysis
Matrix rank analysis (MRA) is a useful tool in identifying the number of independently absorbing
species in complex mixtures91-94. The methods and goals of factor analysis are similar to that of
85
MRA95. Thus both of these statistical methods are discussed in this section, and the differences are
pointed out.
It has been shown, that the rank of the experimental absorption matrix gives the number of
independently absorbing species if Beer’s Law is valid92, 94. The experimental absorption matrix (A)
is given by Equation 43.
(43)
where j
number of wavelengths
i
number of measurements
p
number of independently absorbing species
Ar,s
absorbance measured at wavelength s in measurement r, normalized to unit path length
Cr,t
concentration of species t in measurement r
et,s
molar absorptivity of species t at wavelength s
C
concentration matrix
e
molar absorptivity matrix
The rank of A is the smaller of the number of linearly independent rows or columns96. It is shown
that the rank of A is
rank (A) = min {rank (C), rank (e)} # min{i, j, p}
86
(44)
If the number of measurements and the number of wavelengths are higher than the number of
absorbing species:
rank (A) # p
(45)
Thus, in theory, the number of independently absorbing species could be determined by
determining the rank of A. However, in real measurements, measurement errors are introduced into
the absorbance measured, resulting in significantly higher rank than the number of absorbing species.
Fortunately, there are methods which can be used for determining the number of independently
absorbing species, despite the measurement errors. One of these methods is outlined below.
Matrix A can be decomposed into so-called eigenvectors (X) and their corresponding
eigenvalues (l). The relationship between A, X, and l is defined by Equation 46. The eigenvalues
give a measure of the importance of the corresponding eigenvector.
AX = lX
(46)
If the eigenvectors are sorted according to their eigenvalues from the highest to the lowest, the
first eigenvector contains the most information about A and the last eigenvector contains the least
information about A. The first p eigenvectors contain the information which corresponds to the
independently absorbing species and the rest of the eigenvectors correspond to the measurement
errors. It means that the first p eigenvectors describe A within experimental error. If the number of
selected factors is lower than p, the experimental data would not be reproduced accurately. If the
number of selected factors is greater than p, some of the experimental errors would be included.
Thus, determining p is important.
One method for determining p is to calculate the standard error of each eigenvalue based on an
initial estimate for the error associated with the measurement. This initial error estimate can be the
87
standard deviation of a spectrophotometric measurement. In this way, p is determined by the number
of non-zero eigenvalues. However, this method is greatly sensitive to the accuracy of the initial
estimate of the error.
On the other hand, this method gives a good possibility to filter erroneous data from large data
sets95. Filtering is accomplished by excluding rows or columns from the calculations, one by one. If
the excluded row or column does not contain erroneous data, the eigenvalues and their associated
standard errors are not changed significantly, thus the number of non-zero eigenvalues remains the
same. If the excluded row or column contains errors, the eigenvalues and/or their associated standard
errors change significantly, resulting in a change in the number of non-zero eigenvalues.
The other method to determine p is to calculate the so-called residual absorbance curves. Here
n absorbing species is assumed. First, the eigenvalues and eigenvectors are determined and sorted by
their value from the largest to the lowest. The first n eigenvectors, corresponding to the first n
species, are excluded from A. The resulting matrix gives the residual absorbance curve.
By plotting the residual absorbance curves, it is possible to decide whether n equals p or not. If
the residual curve shows a systematic deviation, it is a good indication that n < p. However,
comparing the residual curves for n and n + 1 absorbing species is important. If the residual curves
in the two cases are not significantly different, the number of absorbing species is n.
It is important to note that the determined eigenvalues or eigenvectors do not have real, physical
meaning. These eigenvectors are only abstract factors which give insight to the complexity of the
chemical system.
In the current research, matrix rank analysis was performed by using the program of Gábor
Peintler97. The use of this program is demonstrated in an article by Peintler et al 95.
88
The program is controlled by using a configuration file. This file is used to switch on/off the
various procedures, give the name of the input file, etc. Through this configuration file, the MRA
program can calculate the eigenvalues and their standard errors, exclude temporarily single or
multiple rows or columns (without actually removing them from the input matrix), and calculate
residual curves.
3.2.2. Determination of formation constants
The formation of various complex species is important in many fields. For example, complexes
are important in preparing antitumor drugs98, MRI contrast agents99, and analytical reagents100-102. For
this reason, the details of the complex formation, the determination of the formation constants and
the molar absorptivities of the complexes are discussed in detail in several papers96, 103-105. Thus, only
a short overview of these methods is given here, with specific reference to the use of computer
programs for determining formation constants and molar absorptivities.
The first step in the determination of the formation constants is to understand the chemical
system. This understanding means to know what chemical reactions take place (complexation and
protonation reactions) and to know the composition of the complexes that are present in the system.
The composition of the complex is most conveniently determined by using Job’s method88. In this
method, the concentration ratio of the two components is varied and a characteristic physical
parameter (e.g., absorbance) is measured. The composition of the complex is given by the extreme
(either minimum or maximum) of the absorbance vs. ratio plot.
Once the basic understanding of the system is established, the calculation of formation constants
and molar absorptivities can be done. The starting point of such calculations is the following mass
balance equation106.
89
(47)
where k
number of components
n
total number of species
Ci
total concentration of the components
Sj
species j in the system
cj
equilibrium concentration of species j
ajk
stoichiometric number
bj
formation constant of species j
The use of these equations requires initial estimates for the formation constants. Based on these
values, the equilibrium concentrations for each of the species are determined by using various
algorithms (e.g., Newton-Raphson) to solve the equation system107.
By using the known molar absorptivities of the components, the total concentrations of the
components, and the absorbance measured, the unknown molar absorptivities of the complexes are
determined. By using this molar absorptivity, the absorbance is calculated (Acalc) and the formation
constant is refined until a minimum of (Ameas–Acalc)2 is reached.
Currently, several programs are available for the determination of formation constants of
complexes either based on spectrophotometric or potentiometric data106. Many of these programs are
based on the previously described procedure. Some differences between the programs include the
algorithms which are used to solve Equation 47 and to minimize the square of the error between the
90
calculated and measured absorbance. Another difference is whether they are general or specific
programs. Specific programs are written to determine the formation constants in a specific system
and in general, some form of simplification108 is used. On the other hand, general programs are
written in a way, that the information about the system (number and composition of complexes) can
be entered along with the experimental data.
3.2.3. PSEQUAD
PSEQUAD109 (Potentiometric and/or Spectrophotometric Equilibrium Data Using Analytical
Derivatives) is a general program for determining the formation constants and molar absorptivities
of complexes106. The program has been written in Fortran and can perform the calculations based on
spectrophotometric and/or potentiometric data. When both spectrophotometric and potentiometric
data are given, they are handled separately. The program takes the input in a rigidly arranged form
from a control file. The lines in this control file turn on/off functions of the program, give options,
etc. The control file contains the input data as well.
As part of the input, the composition of the system is defined through a composition matrix. The
composition matrix has the following general form.
(48)
where Si
ak,l
species i in the system
stoichiometric number
91
The first k rows contain the components and they form a unit matrix. The second part of the
matrix contains the composition information about the complexes. The stoichiometric numbers are
given as positive integer values. In case of self-dissociation of a component (e.g., water) only one of
the products is considered as a component and the other is entered as a complex with a negative
stoichiometric number. For example in the case of water only H+ (OH–) can be considered as a
component. The other, OH–, has a stoichiometric number of –1. The same is true if one of the species
has a deficit in one of the components. For example, a metal hydroxide complex (MOHx+) has a
stoichiometric number of –1 with respect to H+.
In addition to the composition matrix, the program requires the base ten logarithm of the
formation constants for all species. In case of the components, the logarithm of the formation
constants is 0.0. For the complexes with unknown formation constants, an initial estimate is needed.
When spectrophotometric data is analyzed, the program requires the concentration of the components
and the measured absorbances for each measurement. Furthermore, the known molar absorptivities
are required.
The output of the program contains the initial parameters (to make it easier to check if correct
values were used), the refined parameters in each calculation step, the value of the fitting parameter,
the concentration of each species in each measurement, and the residuals for all data points. The
fitting parameter is defined with the following equation.
(49)
where f
n
fitting parameter
number of spectra
92
Wi
weighing factor
Ri
residual of the measurement (Acalculated –Ameasured)
D
degrees of freedom
Despite the advantages of its use, the program has some limitations. One arises from the rigid
input structure. In this input structure only two digits are available for the number of experiments and
wavelengths, thus the highest number of spectra and wavelengths which can be analyzed with this
program is 100. Furthermore, the program is unable to perform calculations when truncated spectra
are present (in which, at some wavelengths, the absorbance is higher than 2). This latter is a severe
limiting factor in this research, as the large excess of chlorite ion in some of the experiments would
severely limit the wavelength range in which the molar absorptivity can be determined.
3.2.4. Excel workbook for the determination of formation constants
Excel provides a convenient alternative to the currently existing programs for the determination
of formation constant. Excel can perform iterations, which makes it simple to solve the mass balance
equation (Equation 47). A useful programming language, Visual Basic for Applications (VBA) is part
of the software which can be easily used for automating repetitive tasks. These two features make
Excel a good alternative.
In this case, a Visual Basic program (macro) is written, which calculates the concentration of the
species in the system by using a predefined set of formation constants. These concentrations and
absorbances measured are used to determine the molar absorptivity of the complex. The sum of the
square of the residuals is calculated and is plotted against the logarithm of the formation constant.
Based on this plot, the value of the formation constant can be determined by finding the minimum.
93
The advantage of this method is that it can handle even truncated spectra. Furthermore, the
number of wavelengths and experiments is much less limited (the maximum number of rows and
columns in Excel are 65536 and 256, respectively).
3.3. Experimental
Preparations of several reagents are described in Chapter 2. These procedures are not repeated
here. This section gives only the procedures which are not detailed in the previous Experimental
section.
3.3.1. Purification of sodium chlorite
Sodium chlorite is a very reactive compound. Therefore, it is not commercially available in pure
form. The approximate composition of the available solid can be found in an American Waterworks
Association (AWWA) standard21. Solid NaClO2 is composed of about 80% sodium chlorite and
various impurities, mainly carbonate and chloride ions.
To use this sodium chlorite for accurate research, sodium chlorite needs to be purified. This was
accomplished by using a previously published method26,
110
. The purification was performed as
follows:
1.
An alcoholic suspension was prepared from the technical grade sodium chlorite, stirred for
15-20 minutes and filtered.
2.
From the filtered solid, a concentrated solution was prepared by using about 40 °C water. The
water was added slowly to make sure that the solution is very concentrated.
3.
In case of any precipitate formation, it was removed by vacuum filtration. Saturated
Ba(ClO4)2 was added to the solution until precipitation occured. This step removes carbonate
ion, which is one of the major impurities in technical grade sodium chlorite. To make it easier
94
to see the precipitate formation, filtering the solution was necessary from time to time when
it became cloudy.
4.
When no more precipitate was formed, the precipitate was removed by vacuum filtration. The
filtrate was checked that no excess barium or carbonate ions were present in the solution. This
was accomplished by adding saturated Ba(ClO4)2 or saturated Na2SO4 solutions to the filtrate.
5.
The solution was mixed with about 2 L of acetone in a large beaker. The mixture was cooled
in a dry ice – ethanol cooling bath for ~ 30 minutes.
6.
The precipitate was filtered and dried by using vacuum filtration. The solid was dissolved in
~40 °C water to make a saturated solution.
7.
Steps 5 and 6 were repeated twice to improve further the purity of the sodium chlorite.
8.
The purified sodium chlorite was placed in a vacuum desiccator and dried for about a week
over anhydrous P2O5. The P2O5 was replaced when it became saturated with water.
The purified sodium chlorite was stored in a vacuum desiccator over P2O5. The purity of the
obtained NaClO2 was determined by iodometric titration9, 30. The purity of the purified solid sodium
chlorite was always 99+ % (m/m).
3.3.2. Preparation of sodium perchlorate solution
For ionic strength adjustments, concentrated (5-7 M) sodium perchlorate solution was used.
Commercially available sodium perchlorate contains sodium chlorate. Therefore, concentrated sodium
perchlorate solutions were prepared by neutralizing sodium carbonate with concentrated perchloric
acid111.
1.
Sodium carbonate was neutralized in a large beaker by dropwise addition of concentrated
(about 85%) perchloric acid. Due to the large neutralization heat, the mixture was cooled in
an ice bath.
2.
After adjusting the pH of the solution to ~9-10, it was allowed to stand for at least 3-4 days.
At the beginning the pH was checked frequently.
95
3.
The solution was filtered on a glass filter. The pH of the filtrate was adjusted to 7-8 and the
solution was allowed to stand again for at least 3-4 days.
4.
Following another filtration, the pH of the solution was adjusted to 2 and the solution was
boiled for a day to remove the CO2 dissolved. The evaporating water was replenished by
TDW. However, the water was carefully added to keep the solution concentrated.
5.
The solution was allowed to cool and checked that no crystals were present. The pH was
adjusted to 7 with carbonate free sodium hydroxide.
6.
This solution was concentrated by boiling so that at room temperature only a small amount
of NaClO4 would precipitate. The solid was filtered and discarded.
7.
The mother liquor was concentrated by boiling and precipitating the sodium perchlorate in
an ice bath. The solid, which is high purity sodium perchlorate, was filtered.
8.
Concentrated solutions (~6-7 M) were prepared from this solid. If these solutions had
significant absorbance above 235 nm in a 1 cm cell, steps 6-8 were repeated. Any absorbance
above 235 nm would indicate the presence of contaminants, e.g., nitrate ion.
The concentration of the sodium perchlorate solution was determined by drying a known volume
at 120°C and weighing the mass of the solid sodium perchlorate.
3.3.3. Determination of the molar absorptivity of ClO 2 and chlorite ion
The molar absorptivities of ClO2 and chlorite ion were determined for both a spectrophotometer
(Agilent 8453) and a stopped-flow instrument (Applied Photophysics SX-18MV). The concentrations
of ClO2 stock solutions were determined prior to the absorbance measurements. Chlorine dioxide
solutions were titrated according to the previously described iodometric procedure (Chapter 2). The
stock solution was diluted to prepare the standard ClO2 solutions. The diluted solutions were
transferred into a quartz cuvette or into the syringes of the stopped-flow instrument(SF) and the
absorbance of the solution was measured in triplicate. Chlorine dioxide concentrations were varied
96
between 3.11×10–4 M and 4.70×10–3 M for the spectrophotometric measurements, and the
concentration range for the SF calibration was from 3.40×10–4 M to 3.40×10–3 M.
Chlorite ion stock solutions were prepared by dissolving purified sodium chlorite in TDW. The
concentration of chlorite ion was measured by the previously described iodometric procedure
(Chapter 2). The stock solution was diluted to prepare the standard chlorite ion solutions. The
absorbance of the solutions was determined in triplicate. The concentration was varied between
3.25×10–3 M and 1.76 M for the spectrophotometric measurements, and between 0.10 M and 1.99 M
for the SF measurements. The absorbance spectra of ClO2 and chlorite ion are shown in Figure 22.
Figure 22. The molar absorptivities of ClO2 and chlorite
ion. The inset shows the 320 nm to 450 nm region. —
Sodium chlorite, — ClO2
97
3.4. Problems with the spectrophotometric measurement of the Cl2O4–
complex
The spectrophotometric measurement of the Cl2O4– complex is problematic. The problems are
summarized below.
!
Chlorine dioxide is volatile and unstable: this means that when measuring the absorbance of
ClO2 solutions, care must be taken to avoid its loss either through evaporation or
decomposition reactions.
!
Keq and e are not independent variables: numerical methods are necessary, which are able to
use large amount of data for determining these values.
!
The spectra of the complex and ClO2 overlap.
!
Small absorbance change: the molar absorptivities of the complex and ClO2 differ significantly
only at longer wavelengths. Thus, to get measurable absorbance changes, a relatively high
concentration of the complex is necessary. This requires concentrated chlorite or ClO2
solutions. If concentrated ClO2 solution is used, the wavelength range of the measurements
is limited due to the high molar absorptivity of ClO2. Therefore, it is necessary to use
concentrated chlorite ion solutions.
3.5. Long-period grating (LPG) sensor results
Due to the limitations of the spectrophotometric measurements, alternative methods were
explored. One such method is to measure the refractive index of the solution. This was accomplished
by using an LPG sensor. In this research, a Luna Innovations FiberPro USB unit was used with bare
LPG fibers from Luna Innovations, which operated around 810-830 nm (the position of the peak in
air).
98
Fiber gratings are periodic perturbations of the refractive index of an optical fiber112. In the case
of long-period grating, the period is generally on the order of 100 µm to 1 mm112. The LPG sensor
consists of a central optical grade fiber and a surrounding cladding112-114. If light is passed through the
fiber, the resulting spectrum and the center wavelengths of the bands are dependent on the light
source, on the period of the LPG, and on the surrounding environment. Changes in these parameters
result in a change of the central wavelength of the bands.
LPG sensors can be used effectively for measuring the change of the refractive index of the
surrounding solution112, 115. The response of the sensor to refractive index change is not linear. The
higher limit of the refractive index, which can be measured by an LPG sensor, is dependent on the
refractive index of the cladding, because it needs to be higher than the surrounding environment. Bare
LPG fibers are non-selective112, but selectivity can be achieved by coating the fiber with a selective
sensing element112. For example, in this way it is possible to create biological sensors116 or a sensor
for measuring copper113.
The signal of the LPG sensor is a single peak in air (Figure 23). When the sensor is immersed
into a solution, two peaks are observed (Figure 24). The appearance of these two peaks instead of
one, is due to the loss of light, which exits the fiber, creating a hole (or valley) at this position114. The
wavelength of this valley is the measured value in these measurements. The position of this valley is
dependent on the refractive index of the surrounding solution. The wavelengths of the valley in
solutions and the peak in air are not reproducible on different fibers, due to manufacturing
imperfections114.
99
First, repeatability, reproducibility, and precision of the response of a sensor in water were
determined. Precision was determined by using one fiber (#15) and twenty replicate measurements.
The average value for the wavelength of the valley was 820.5 nm and the standard deviation was 0.29.
Figure 23. Signal of an LPG sensor in air.
Figure 24. Signal of an LPG sensor in water
100
Repeatability was tested for one fiber (#15) on a single day and on three different days. The
results are summarized in Table 16. As these results reveal, significant variation is observed in the
wavelength of the valley, and the standard deviation of this wavelength is high.
Table 16. Repeatability of the wavelength of the valley in water
for sensor #15.
Day/Replicate
Average
Standard
Deviation
# of
replicates
Day 1/Replicate 1
822.1
0.19
5
Day 1/Replicate 2
822.9
0.18
5
Day 1/Replicate 3
823.1
0.20
5
Day 1/Replicate 4
822.6
0.23
5
Day 1/Replicate 5
822.7
0.24
5
Day 2/Replicate 1
820.5
0.29
20
Day 3/Replicate 1
822.1
0.19
5
Reproducibility was tested by using three different sensors (#3, #15, and #16) to determine the
wavelength of the valley in water. The results of this measurement are summarized in Table 17. As
it can be seen from the data, the wavelength of the valley is significantly different for various fibers.
If the produced sensors were similar to each other, the wavelength of the valley would not be
significantly different for the various fibers.
Table 17. Comparison of the position of the valley for
three sensors.
Fiber
Average
Standard
Deviation
# of
replicates
#3
830.6
0.13
5
#15
822.1
0.19
5
#16
817.3
0.13
5
101
Because of these problems with repeatability and reproducibility in the valley position in water,
it is necessary to use the difference in the position of the valley during subsequent work. The
difference in the valley position can be given with the following equation.
ldiff. = lsample – lwater
(50)
3.5.1. The calibration of the LPG sensor for the determination of chlorite ion
concentration
Due to the repeatability and reproducibility problems in water, accuracy, repeatability, and
reproducibility of the response of the sensor were tested with sodium chlorite solutions as well.
The accuracy of the response of the LPG sensor was tested with seven different NaClO2
concentrations between 0.50 M and 2.98 M. In these measurements twenty replicates were used. The
results are summarized in Table 18 and the calibration curve is shown in Figure 25.
Table 18. Results of the calibration of LPG sensor for chlorite
ion.
c(NaClO2)
lvalley (nm)
Standard
deviation
ldiff.
0
822.4
0.29
N/A
0.5
822.8
0.15
0.32
0.99
821.7
0.14
-0.73
1.24
821.2
0.24
-1.26
1.49
820.5
0.23
-1.97
1.99
819.7
0.26
-2.68
2.49
817.2
0.37
-5.25
102
Figure 25. Calibration curve of an LPG sensor for chlorite
ion. The equation of the line is lvalley = 2.95×[NaClO2] +
824.61, R2=0.965
The precision of the measurement is good and is about the same as in the case of water. The
calibration curve is good, showing good linear response. At the time of the experiments, the effect
of ionic strength on the observed signal was not clear. However, recent results in the laboratory of
Dr. Gilbert Pacey indicate that the change of the signal is dependent on ionic strength changes. Thus,
the observed linear change is due to ionic strength changes and not chlorite ion concentration change
itself.
Calibration curves were created on different days by using the same sensor (#15) and compared.
The results are shown in Figure 26 and Table 19. These results show that the calibration curves are
significantly different on various days. This means that it is necessary to calibrate the sensor daily.
Table 19. Comparison of the calibration curves for
the same sensor on different days.
Day
Slope
Intercept
R2
1
-2.82
2.14
0.928
2
-2.33
0.35
0.941
3
-1.49
-0.3
0.849
103
Figure 26. Calibration curves of an LPG sensor for chlorite ion using
the same fiber on different days. For the parameters of the calibration
curves see Table 19.  Day 1, Ž day 2, • day 3
Calibration curves were constructed on different fibers as well. The results are shown in Figure
27 and Table 20. In these measurements the same chlorite ion solutions were determined by using
three different sensors (#3, #15, and #16) on two days. Each solution was measured five times. As
it can be seen, the calibration curves of the three sensors are significantly different.
Table 20. Parameters of calibration curves for
chlorite ion determination for various sensors
Fiber
Slope
Intercept
R2
#3
-1.74
-0.57
0.958
#15
-2.13
0
0.951
#16
-1.41
1.03
0.971
104
Figure 27. Calibration curves of various LPG sensors using the same
chlorite ion solutions and different fibers. For the parameters of the
calibration curves see Table 20.  #3, Ž #15, • #16
The results of these experiments show that even though the LPG sensor is able to detect chlorite
ion, the response of the sensors for chlorite ion does not show good reproducibility or repeatability.
Furthermore, the currently used bare LPG sensors do not have the required sensitivity to detect small
changes due to the Cl2O4– complex formation. The LPG sensors may become a good alternative for
the measurement of chlorite ions if a suitable coating is found, which would improve the
characteristics of the bare LPG fiber and make it selective for the measurement of chlorite ion.
Furthermore, based on the calibration curve in Figure 25 and the standard deviation of the
measurements, it can be concluded that lowest chlorite ion concentration change that can be detected
is about 0.3 M. However, the concentration change due to the Cl2O4– complex formation is about
10–3-10–4 M. Thus, in its current form the LPG sensor does not have the required sensitivity to detect
this concentration change.
105
3.5.2. The calibration of the LPG sensor for the determination of ClO 2
concentration
Long period grating sensors were tested to determine whether they are able to sense ClO2 and
whether the response of the sensor can be used to measure ClO2 concentrations. Chlorine dioxide
concentrations were in the low concentration range (from 0.02 M to 0.04 M).
The calibration curve is shown in Figure 28. The Figure reveals that no trend was observed in
the valley position as a function of ClO2 concentration. This means that the bare LPG sensor is unable
to determine ClO2 at this low concentration range. However, the use of higher ClO2 is limited for
several reasons. The maximum ClO2 concentration that can be generated by using the current
generation method is about 0.06 M. Furthermore, at medium to high concentration range, a
significant amount of ClO2 is lost through evaporation. Thus, the bare LPG sensor is not useful in
measuring the ClO2 concentrations which are used in this research.
Figure 28 Calibration curve of an LPG sensor for ClO2.
106
3.5.3. The response of the LPG sensor in mixtures of ClO 2 and chlorite ion
The response of a sensor was tested in mixtures of ClO2 and chlorite ion. The valley position is
plotted as a function of Cl2O4– complex concentration in Figure 29. This Cl2O4– complex
concentration was calculated by using the previously determined value of the formation constant
(1.6 M–1). As the Figure indicates, no clear trend is observed in the response of the LPG sensor as
a function of Cl2O4– complex concentration.
Based on these results, it is clear that using bare fibers, the LPG sensor does not have sufficient
sensitivity for the determination of the complex. However, by modifying the fibers with various
chemicals, the measurement of ClO2 and chlorite ion could possibly be improved.
Figure 29. Response of an LPG sensor at different Cl2O4– complex
concentrations. The concentration of the complex was determined
by using Keq = 1.6 M–1.
107
3.6. Initial spectrophotometric results
Initial experiments were performed by using a tandem cell, which is shown in Figure 30. This cell
is primarily designed for kinetic measurements and consists of two compartments that are separated
by a polished quartz window. An opening at the top of this wall allows the mixing of the two reagents
to start the reaction but prevents unintentional mixing. By using this cell, a comparison of the spectra
of chlorite ion and ClO2 can be obtained before and after mixing. The advantage of this cell is that the
solutions can be mixed rapidly without significant loss of ClO2.
Figure 31 and Table 21 show the results of a typical experiment. The Figure shows that the
difference between the unmixed and mixed solutions around 400 nm is nominal. However, this
difference increases with increasing wavelengths. This means that in concentrated ClO2 solutions,
where the absorbance is measured at high wavelengths (450 nm or even 500 nm), the absorbance
increase becomes significant.
a)
b)
Figure 30. a)Photograph and b) schematic drawing of tandem cell
108
If in the measurement – shown in Figure 31 –, the ClO2 concentration had been determined by
using spectrophotometric measurement at 450 nm, the calculated ClO2 would be 3.94×10–3 M
(266 mg/L). This calculated concentration is about 58% higher than the concentration determined by
spectrophotometric measurement at the same wavelength before mixing. As a result, in water
treatment plant applications significantly less ClO2 would be used than required, potentially resulting
in adequate disinfection.
According to the results in Table 21, it can be seen that the error in the ClO2 measurement
increases with increasing chlorite ion concentration. In contrast, at the constant chlorite ion
concentration, the error is lower at higher ClO2 concentration.
Figure 31. Absorbance change of a ClO2 and chlorite ion
mixture before (—) and after mixing (þ). c(ClO2) =
168.4 mg/L (2.5×10–3 M), c(NaClO2) = 112.0 g/l (1.66 M),
A450 nm = 0.130 before mixing, A450 nm = 0.184 after mixing,
path length = 2×0.437 cm
109
Table 21. The effect of the Cl2O4– complex on the spectrophotometric
measurement of ClO2. c(ClO2)calculated = A450 nm/e450 nm(ClO2), % error =
[c(ClO2)calculated–c(ClO2)]/c(ClO2)×100
c(ClO2)
c(NaClO2)
A450 nm
c(ClO2)calculated
% Error
165 mg/L
(2.44×10–3 M)
8.50 g/L
(0.126 M)
0.104
167 mg/L
(2.48×10–3 M)
1.31
165 mg/L
(2.44×10–3 M)
33.9 g/L
(0.502 M)
0.133
213 mg/L
(3.16×10–3 M)
29.1
165 mg/L
(2.44×10–3 M)
76.2 g/L
(1.13 M)
0.162
259 mg/L
(3.85×10–3 M)
57.4
165 mg/L
(2.44×10–3 M)
102 g/L
(1.51 M)
0.169
271 mg/L
(4.01×10–3 M)
64.1
168 mg/L
(2.50×10–3 M)
138 g/l
(2.05 M)
0.181
290 mg/L
(4.30×10–3 M)
72.2
69.0 mg/L
(1.02×10–3 M)
33.1 g/L
(0.491 M)
0.067
107 mg/L
(1.59×10–3 M)
55.7
The results of the initial measurements were used to determine the number of independently
absorbing species. For this, the program97 MRA was used. The number of absorbing species is given
by the number of non-zero eigenvalues. However, it is known that the error associated with the
measurement may present a non-zero eigenvalue. Thus, the number of absorbing species, which is
determined this way, is the maximum number which may be present in a given system and may include
experimental errors.
The program MRA presents another method for determining the number of absorbing species95.
This can be done by determining the residual spectra. The residual spectrum is not a real absorbance
spectrum, it is derived from the measurement matrix. The residual spectrum is determined by
assuming a theoretical number of absorbing species, and their contribution to the absorbance is
subtracted from the original spectrum. The number of absorbing species can be determined by visual
110
inspection of the residual curve (see the Theoretical section of this chapter). If the assumed number
of absorbing species matches the number of absorbing species which are present, the residual
spectrum contains data related only to various errors in the spectrophotometric measurement. Thus,
the resulting residual spectrum becomes featureless and noise-like.
The determined eigenvalues and their corresponding standard errors are presented in Table 22.
From these results it can be seen that the number of non-zero eigenvalues is four. This is higher than
the expected three. The three expected absorbing species are ClO2, chlorite ion, and the Cl2O4–
complex.
Table 22. Results of an MRA run.
Number of spectra = 78, wavelength
range: 395–600 nm
#
Eigenvalue
Standard
error
1
2.024
0.010
2
-0.077
0.013
3
-0.040
0.014
4
0.030
0.021
5
0.027
0.030
The determined residual spectra for 1 to 4 absorbing species are presented in Figure 32. As it can
be seen from this figure, after subtracting the spectra of one and two absorbing species, the residual
spectrum still shows characteristic deviation. However, subtracting the spectrum of the third
absorbing species, the deviation is generally less than 0.01 that is the standard deviation of the
spectrophotometer used . Furthermore, this residual curve does not change significantly when 4 or
5 absorbing species are assumed (residual curve shown only for 4 absorbing species), indicating the
111
low contribution of the fourth and further species to the overall absorbance. Thus, it can be concluded
that only three absorbing species (ClO2, ClO2–, and Cl2O4–) are present in the ClO2–chlorite ion
system.
Figure 32 Residual spectra after assuming — 1, — 2, — 3, and — 4
absorbing species.
3.7. Main spectrophotometric study
The main study was performed by using an Applied Photophysics SX-18MV stopped-flow
spectrophotometer with a photodiode array detector. The temperature was 25°C, the ClO2
concentration was varied between 3.34×10–4 M and 1.67×10–3 M, and the chlorite ion concentration
was varied between 0.30 M and 0.98 M. A constant ionic strength of 1.0 M was maintained by using
NaClO4. A boosted deuterium lamp was used in the measurements, which was directly mounted on
112
the spectrophotometric cell block. This lamp provides sufficiently intensive light from the UV region
up to 700 nm.
During the stopped-flow measurements, one syringe of the stopped-flow was filled with ClO2
solution and the other syringe with chlorite ion solution. In each measurement 100 spectra were
collected (the minimum number of spectra that can be selected) for a total of 400 ms. This
corresponds to 4 ms integration time for each spectrum. Longer integration time would be beneficial
to improve the signal-to-noise ratio. However, this is limited due to the following. Because of the
short measurement times in SF measurements, there is no optical element that opens/closes the slit.
Thus, the light from the lamp continuously illuminates the sample. Due to this uninterrupted
illumination, the photodecomposition of chlorite ion becomes observable after about 0.5 s and
becomes significant after about 1 s. This limits the maximum integration time that can be used.
A sufficiently high number of spectra (more than 200) was collected, and the data were analyzed
by using the program109 PSEQUAD and the previously described Excel workbook. Figure 33 shows
the PSEQUAD results and Figure 34 shows the Excel results. The two data sets, which were used
for fitting, were different in the used wavelength range because PSEQUAD can fit the data only when
full rows and columns are present. This means that for PSEQUAD fitting a smaller wavelength range
was used. The calculated formation constants are formal formation constants because their values
were determined on the basis of the concentrations of the species.
113
Figure 33. Results of fitting of stopped-flow data by
using PSEQUAD. Temperature = 25 °C
The parameters which were minimized in these calculations are:
fitting parameter in PSEQUAD calculations:
(14)
where f
fitting parameter
n
number of spectra
Wi
weighing factor
Ri
residual of the measurement (Acalculated–Ameasured)
D
degrees of freedom
error2 in case of the Excel workbook
(51)
114
Figure 34. Results of fitting of stopped-flow data by
using Excel worksheet. Temperature = 25 °C
The comparison of the two figures shows that the two methods give similar results. However,
when PSEQUAD was used to fit the data, the fitting parameter does not change within a range of log
Keq (from 0.4 to 0.7). These values correspond to the range of Keq from 2.5 to 5.0. This means that
by using PSEQUAD it is not possible determine the exact value of Keq.
Based on the Excel worksheet calculations, the formation constant of the complex is 5.0 M–1.
By using this value of the formation constant, the molar absorptivity of the Cl2O4– complex was
determined as a function of wavelength. This molar absorptivity is compared with the molar
absorptivity of ClO2 and chlorite ion in Figure 35 and Table 23.
The results show that at 470 nm and higher, where ClO2 has very low absorptivity (practically
zero), the complex still has significant absorbance. Thus, if the absorbance of a ClO2 solution is
measured at this high wavelength and the solution contains significant amounts of chlorite ion, the
measured absorbance would be mainly due to the complex. This would result in false high ClO2
concentration readings.
115
Figure 35. Comparison of the molar absorptivity of the
various species in the chlorite ion–ClO2 system.
The absorbance change at a given wavelength is due to the decrease in the ClO2 concentration
and the increase in the complex concentration. The absorbance change from chlorite ion is neglected
because its molar absorptivity is low and its concentration change is minimal (chlorite ion is assumed
to be in excess).
(52)
where DA
absorbance change
c(ClO2)
total concentration of ClO2
[ClO2]
equilibrium concentration of ClO2
[Cl2O4–] = c(ClO2) – [ClO2]
equilibrium concentration of the Cl2O4– complex
116
Table 23. Molar absorptivity of ClO2, NaClO2, and Cl2O4–
Wavelength
(nm)
Molar absorptivity (M–1cm–1)
ClO2
NaClO2
Cl2O4–
400
579.8
0.05
544.6
410
414.0
0.03
391.0
420
273.1
0.02
271.2
430
158.1
0.01
172.3
440
82.6
0.01
111.6
450
45.6
0
75.3
460
22.0
0
52.2
470
9.3
0
39.9
480
4.7
0
33.5
490
2.1
0
30.1
500
1.2
0
25.6
525
0.2
0
17.9
550
0
0
13.4
At constant ClO2 and chlorite ion concentration, the concentration of the Cl2O4– complex is
constant. In this case, the absorbance change due to the complex is dependent on the difference
between the molar absorptivities of ClO2 and the complex. By using the molar absorptivities of the
two species, one can see that the below ~430 nm, the absorbance is decreasing and above this value
the absorbance is increasing due to the presence of the complex. It can be concluded that with
increasing wavelengths, the absorbance increase is becoming larger due to the increase in the molar
absorptivity difference.
117
3.7.1. The effect of temperature on the equilibrium
In practical applications, the temperature of a ClO2 solution may be significantly different from
room temperature. In such cases, the value of the formation constant is expected to be different from
the value determined at room temperature. To be able to correct for the presence of the Cl2O4–
complex at these temperatures, determining the effect of the temperature on the formation constant
is necessary.
These measurements were taken on the stopped-flow spectrophotometer. The temperature range
was from 15°C to 45°C. The ClO2 concentration was varied from 2.01×10–4 M to 1.67×10–3 M. The
chlorite ion concentration was varied between 0.295 M and 0.984 M. The ionic strength was adjusted
to 1.0 M with NaClO4.
The values of the formation constant at the different temperatures were determined by using the
previously described Excel worksheet and are summarized in Table 24 and Figure 36. In Figure 36,
the logarithm of the equilibrium constant is plotted against the reciprocal of the absolute temperature
to determine the enthalpy of the reaction according to the following equation. The determined value
is DH0 = –43.9 ± 6.5 kJ/mol.
(53)
118
Figure 36. The change of the formation constant of the
Cl2O4– complex with temperature. The equation of the least
squares fitted line is log Keq = 2293.8×T–1 – 7.06, R2 =
0.958
Table 24. The change of the equilibrium
constant with temperature.
Temperature
log Keq
15 °C
0.9
25 °C
0.7
35 °C
0.3
45 °C
0.2
3.8. The structure of the complex
Electron resonance spectra (EPR) can be obtained for ClO2 dissolved in aqueous solutions
because it is a free radical. In concentrated solutions a broad peak is observed that can be resolved
into a four-component hyperfine pattern at low concentrations117. The hyperfine pattern is due to the
interaction of the unpaired electron with the chlorine nucleus (I = 3/2). A typical EPR spectra of ClO2
119
is shown in Figure 37. The average separation of the peaks is 1.68×10–3 T (16.8 gauss), and the
average width is 7.65×10–4 T (7.65 gauss). These values are in good agreement with the previously
reported values of 17 gauss of peak separation and 8 gauss line width118.
Figure 37. EPR spectra of ClO2 and ClO2/chlorite ion mixture at
room temperature. [ClO2] = 2.46×10–3 M in ClO2 solution (—),
[ClO2] = 2.46×10–3 M, [ClO2–] = 5.07×10–2 M in Cl2O4– solution (—).
The EPR spectra were collected with a center field of 3370 G, sweep
width of 200 G, a microwave frequency of 9.424 GHz, modulation
frequency of 100 kHz, modulation amplitude of 10 G, and a power of
0.635 mW.
The Cl2O4– complex contains one unpaired electron and is expected to be EPR active. Figure 37
shows the spectra of a mixture of ClO2 and chlorite ion. Comparison of the EPR spectra of pure ClO2
and ClO2 in the presence of chlorite ion reveals a small shift in the line positions and minor changes
in the line shape. These changes are associated with the formation of the Cl2O4– complex. It was
attempted to create larger changes by increasing the concentration of ClO2 and chlorite ion. However,
the increase in the concentrations resulted in two broad peaks, without the resolved hyperfine pattern.
120
Figure 38 represents the structures which are chemically possible. However, Structure III would
be expected to be relatively unstable and to decompose readily to chlorate ion (ClO3–) and OCl.
Therefore only Structures I, II, and IV were considered in detail.
Figure 38. The possible structures of the Cl2O4– complex.
The following can be deduced based on the EPR spectrum and the most feasible structures. In
Structure I, the two chlorine atoms are directly connected which means that the unpaired electron
would interact with both chlorine nuclei. This would result in a more complex hyperfine pattern,
which is not observed in the EPR spectrum. On this basis Structure I can be ruled out as a possible
structure of the Cl2O4– complex. This agrees with the early findings of Gordon et al 12.
Structures II and IV do not contain direct chlorine-chlorine bonding, thus a similar hyperfine
pattern is expected as in the case of ClO2. For this reason, deciding which is the structure of Cl2O4–
complex is not possible based on the EPR spectrum. However, based on chemical considerations, it
is possible to decide which structure is the most likely. Structure II contains only chlorine-oxygen
121
bonds, whereas in Structure IV two additional oxygen-oxygen bonds are present. These bonds are
expected to be less stable than the chlorine-oxygen bonds. Thus, based on these considerations,
Structure II may be the most likely structure of the Cl2O4– complex.
It is tempting to measure the EPR spectra of ClO2 at low temperatures because of improved
resolution. Improved resolution was observed117 in organic solvents (e.g., CCl3F) at around –110°C.
However, the use of organic solvents has disadvantages in the current work. The solubility of sodium
chlorite is expected to be low in these solvents. Thus, because of the low chlorite ion concentration
and the low value of the formation constant, the concentration of the Cl2O4– complex formed would
not be sufficient to make observable changes.
The use of aqueous solution has its shortcomings as well. At these low temperatures the aqueous
solution would be frozen. According to Ingram118, in the frozen form the lines become “much wider
and less resolved.” This change in the resolution is due to “the anisotropic hyperfine interaction,
which is averaged to zero in the liquid state.” However, in the solid state, the non-zero anisotropic
interaction causes line broadening. Another problem with cooling aqueous solutions of ClO2 is that
at low temperatures (below the boiling point of ClO2) liquid ClO2 may separate from the aqueous
solution. Liquid ClO2 is very unstable and readily decomposes in pure form. Thus, the measurement
of EPR spectra for ClO2 in aqueous solution at low temperatures is not a good option in determining
the exact structure of the Cl2O4– complex.
The Cl2O4– complex is known to undergo a decomposition reaction86. The identification of the
products of this decomposition reaction can provide further evidence in deciding between the possible
structures. However, this decomposition reaction is very slow (total decomposition may take several
days or even weeks) in the chlorite ion and ClO2 concentration range reported here. At higher
122
concentrations the decomposition becomes faster, but those conditions present a severe explosion
hazard.
3.9. Methods to eliminate the interference of the complex
Many ways can be considered for minimizing or eliminating the interference of the Cl2O4complex, but some of these may not be practical or simple enough to be useful in the field. The most
satisfactory of these methods is to eliminate the presence of chlorite ion from the ClO2 solutions. This
could be accomplished by keeping the ClO2 generator carefully under control. However, this may not
be feasible always, thus considering other options is necessary.
The relationship between the measured and “true” ClO2 concentrations is linear if the chlorite
ion concentration is relatively constant. If chlorite ion can not be eliminated from the ClO2 solution,
but can be kept at a relatively constant value, the correction for the complex can be straightforward.
By using approximations in the equations that describe the equilibrium system, the following simple
equation can be derived for the relationship between the measured absorbance and “true” ClO2
concentration:
(54)
where: A8:
,8 :
the measured absorbance at 8 wavelength
the molar absorptivity of the various species at 8 wavelength (these values are given
in Table 23 for ClO2, NaClO2, and Cl2O4-)
123
Table 25 demonstrates the use of Equation 54. The first, c(ClO2) column in this table is the true
ClO2 concentration as determined by iodometric titration. The fourth column, the calculated ClO2
concentration is based on the measured absorbance, which was divided by the molar absorptivity of
ClO2 at the given wavelength to obtain the concentration. The sixth column, corrected ClO2
concentration, is calculated from the measured absorbance by using Equation 54. Comparison of the
% error for the calculated and corrected ClO2 concentrations shows that when Equation 54 is used
the error in the ClO2 concentration is reduced significantly. The error associated with the corrected
ClO2 concentrations is about on the order of the error associated with the spectrophotometric
measurements. Thus, the error resulting from the absorbance correction equation does not
significantly influence the accuracy of the measurement of ClO2.
If the chlorite ion concentration is relatively high and varies within a wide range, using other
methods to correct for the presence of the complex is necessary. There are two good possibilities.
The first is to measure the absorbance of the ClO2 solution at two wavelengths, one being sufficiently
high (between 500 nm and 550 nm) such that only the complex has an absorbance. A second
wavelength is selected (depending on the ClO2 concentration, in the 400-470 nm region) where both
ClO2 and the complex have measurable absorbance. There are commercially available ClO2 analyzers,
which use filters to select the wavelengths (between 400-600 nm) where the absorbance is measured.
These analyzers present a good option for the measurement of the absorbance of the solutions at two
wavelengths.
124
Table 25. The use of Equation 54 to correct for the presence of the Cl2O4– complex. c(ClO2)calculated =
A449.7 nm/e449.7(ClO2), % error = [c(ClO2)calculated - c(ClO2)]/c(ClO2)×100
c(ClO2)true
c(NaClO2)
A449.7 nm
c(ClO2)calculated
%
error
c(ClO2)corrected
% error
113 mg/L
(1.67×10–3 M)
66.4 g/L
(0.984 M)
0.107
167 mg/L
(2.48×10–3 M)
48.7
118 mg/L
(1.76×10–3 M)
5.5
113 mg/L
(1.67×10–3 M)
39.8 g/L
(0.590 M)
0.099
156 mg/L
(2.31×10–3 M)
38.2
115 mg/L
(1.71×10–3 M)
2.4
113 mg/L
(1.67×10–3 M)
26.5 g/L
(0.394 M)
0.088
138 mg/L
(2.05×10–3 M)
22.9
106 mg/L
(1.57×10–3 M)
-5.7
101 mg/L
(1.50×10–3 M)
66.4 g/L
(0.984 M)
0.096
151 mg/L
(2.24×10–3 M)
49
107 mg/L
(1.58×10–3 M)
5.2
101 mg/L
(1.50×10–3 M)
39.8 g/L
(0.590 M)
0.09
142 mg/L
(2.10×10–3 M)
39.7
105 mg/L
(1.55×10–3 M)
3.2
101 mg/L
(1.50×10–3 M)
26.5 g/L
(0.394 M)
0.086
135 mg/L
(2.00×10–3 M)
33.2
104 mg/L
(1.53×10–3 M)
2.1
101 mg/L
(1.50×10–3 M)
19.9 g/L
(0.295 M)
0.074
117 mg/L
(1.73×10–3 M)
15.3
95.3 mg/L
(1.41×10–3 M)
-5.9
78.8 mg/L
(1.17×10–3 M)
66.4 g/L
(0.984 M)
0.076
119 mg/L
(1.77×10–3 M)
51.3
83.0 mg/L
(1.23×10–3 M)
5.3
78.8 mg/L
(1.17×10–3 M)
39.8 g/L
(0.590 M)
0.073
114 mg/L
(1.69×10–3 M)
44.7
83.6 mg/L
(1.24×10–3 M)
6.1
78.8 mg/L
(1.17×10–3 M)
26.5 g/L
(0.394 M)
0.068
107 mg/L
(1.59×10–3 M)
36.2
81.8 mg/L
(1.21×10–3 M)
3.8
78.8 mg/L
(1.17×10–3 M)
19.9 g/L
(0.295 M)
0.06
94.8 mg/L
(1.41×10–3 M)
20.3
77.0 mg/L
(1.14×10–3 M)
-2.3
56.3 mg/L
(8.35×10–4 M)
19.9 g/L
(0.295 M)
0.044
69.6 mg/L
(1.03×10–3 M)
23.7
56.0 mg/L
(8.31×10–4 M)
-0.5
The other method for correction is to measure the absorbance at one wavelength, where both
ClO2 and Cl2O4– complex absorb light (in the 400–470 nm range). The ClO2 is purged from the
sample and the concentration of chlorite ion is determined independently by using other methods,
125
such as amperometric titration. From the chlorite ion concentration and the measured absorbance
value the “true” concentration of ClO2 can be determined by using Equation 54.
This method, however, has several disadvantages. First, it is not a viable method for on-line ClO2
analyzers because it requires a sampling step to determine the chlorite ion concentration. Second, the
purging step, which takes a long time, limits the usability of the method (e.g., if the ClO2
concentration is higher than 500 mg/L, it may take up to one hour). This purging step can lead to
significant error in the measurement of chlorite ion. If the purging is not carried out for sufficiently
long time, some ClO2 remains in the solution, which is measured along with chlorite ion in the
titration. This inaccurate chlorite ion concentration will result in increased inaccuracy in the corrected
ClO2 concentration.
It is necessary to caution against the use of too high wavelengths (e.g., 500 nm and higher) for
the determination of ClO2, because at such high wavelengths ClO2 has practically zero absorbance.
Any absorbance measurement taken at such wavelengths is inherently erroneous, resulting in
inaccurate ClO2 readings.
3.9.1. Comparison of Cl2 O 4 – complex with other similar species
Several intermediate species are reported in oxychlorine chemistry to form in a similar way as
the Cl2O4– complex. Some examples17, 18, 119 are Cl2O3 or Cl2O3–, and120 Cl4O102–. The existence of some
of these species is well-accepted. These species are generally suggested as intermediates in complex
chemical systems where various chemical reactions take place. However, when interpreting such
phenomena it is important to consider carefully the possibilities and choose the species that has the
fewest new, unsupported assumptions121.
126
For example, the species Cl2O3 or Cl2O3– can be considered good alternatives in a reaction
mechanism because the structures can be described based on well-established chemistry. Furthermore,
Cl2O3 is well-known in atmospheric chemistry and its UV spectrum was reported in the gas phase119.
In solution the Cl2O3 complex is predicted to be an intermediate in the hypochlorous acid – ClO2
reaction17, 18. In solution, however, less is known about this species. For example, no UV-visible
spectrum is available.
In contrast, tetrachlorinedodecaoxide120 (Cl4O102–) is an unlikely species, because its proposed
structure contains bonds that are not stable, and their formation is not known in other molecules. A
series of papers postulate Cl4O102– as an important chemical intermediate120. The authors cite
microbiological evidence for its existence. However, no detailed chemical properties have been
reported.
3.10. Conclusions
The spectrophotometric measurement of ClO2 can be severely inaccurate in the presence of high
concentrations of chlorite ion, due to the formation of the Cl2O4– complex. The contribution of the
Cl2O4– complex to the overall absorbance of such a solution can be significantly higher than the
absorbance of ClO2.
By using numerical methods, the formation of Cl2O4– complex in the ClO2–chlorite ion system
has been confirmed. The formation constant of this complex has been determined to be 5.0 M–1.
Based on this value, the molar absorptivity of the complex has been calculated as a function of
wavelength.
127
Based on this formation constant, adjusting the spectrophotometric measurements to reduce the
inaccuracy of the measurement is possible. By using Equation 54, it is possible to reduce the error
of the ClO2 concentration determined to a level that is comparable to the error associated with
spectrophotometric measurements.
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4. Measurement of Reactive Species and Intermediates in
Mixed Disinfectant Solutions: The Dissolved Chlorine
(Free Available Chlorine, FAC)–ClO 2 System
There is a high demand for fast acting, strong disinfectants in various industries, for example in
the medical field or water treatment. The possibilities to improve the efficacy of simple disinfectants
(that contain only one microbiologically active species) are limited. The efficacy generally can be
increased by increasing the concentration of the active species, either directly by increasing its total
concentration in the disinfectant or indirectly by varying the pH of the disinfectant. However, these
changes may lead to other problems, e.g., increased toxicity of the disinfectant or corrosion problems.
Many of these problems can be overcome by using mixed disinfectant solutions.
Synergism between the components of the mixed disinfectant solution can greatly promote the
use of such disinfectants. Synergism is a general term, which can be used when the details or reasons
for the increased microbiological efficacy are not known. It is generally due to the interaction of the
components that form new, microbiologically active species. The interaction between the components
can be an equilibrium process or an irreversible reaction. One of the advantages of the latter case is
that the initial active species are used up in the process. Thus, the remaining disinfectant solution is
less toxic and easier to depose. However, this is a disadvantage of such disinfectants because it limits
the use of the disinfectant and it needs to be used shortly after preparation. The reaction between the
components provides another means to adjust the microbiological activity of the disinfectant solution.
The concentration of the intermediates, which are assumed to be potent disinfectants, changes
characteristically during the reaction. This concentration profile can be changed by adjusting many
129
factors (including temperature and concentration of catalytic or inhibitor species) that also alter the
efficacy of the disinfectant solution.
The hypothesis of this research was that creating more effective and faster acting disinfectant
solutions is possible by mixing chlorine and ClO2 solutions that react with each other. Increased
efficacy is expected for these solutions due to the formation of reactive intermediates in the FACClO2 reaction. In this research it was tested if using kinetic and microbiological results to estimate
the effectiveness of various intermediates formed in this system is possible by using simple methods.
Usually the measurement of such species is complicated and requires detailed measurements.
The objectives were to create mixed disinfectant solutions which are more rapidly acting than
the current disinfectants and characterize the new mixed disinfectants kinetically and
microbiologically. After the initial tests, improving the efficacy of the disinfectants is possible by
combining the results of the microbiological and kinetic experiments.
4.1. Theoretical
4.1.1. The FAC-ClO 2 reaction
Chlorine dioxide and dissolved chlorine react with each other at a relatively fast rate17, 18. During
this reaction several reactive intermediates are formed which may contribute to the disinfection
efficacy of the mixture of the two disinfectants. The moderately fast reaction between dissolved
chlorine and ClO2 is an important factor in selecting this disinfectant couple.
To prepare a fast acting disinfectant, a moderately fast reaction is needed in which the
intermediates are active disinfectants. The reasons for this can be summarized as follows. If the
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reaction takes place too fast, the original disinfectants and the intermediates are present in high
concentration only for a short time. This is not advantageous for the disinfection because the
disinfection process is relatively slow. Thus, the chemical reaction would consume the disinfectants
before significant disinfection would take place. On the other hand, if the reaction is too slow, the
concentrations of the intermediates are always low during the disinfection process and their effect is
negligible. Thus, this last case does not provide improved disinfection.
The FAC-ClO2 reaction has been studied in detail previously17, 18. These studies included kinetic
and mechanistic information. This section gives a summary of these results and how they can be used
in designing a mixed disinfectant system which takes advantage of the FAC-ClO2 reaction.
Below pH 9, the reaction is described with the following equation.
2 ClO2 + HOCl + H2O = 2 ClO3– + Cl– + 3 H+
(55)
At pH values above 9 the decomposition of ClO2 becomes predominant
2 ClO2 + 2 OH– = ClO2– + ClO3– + H2O
(3)
The stoichiometry of the FAC–ClO2 reaction is kinetically controlled and can be given as a
combination of Equations 3 and 55. The detailed mechanism of this reaction is given as18
OCl– + H+ ¾ HOCl
(fast equilibrium)
(56)
ClO2 + OCl– ¾ Cl2O3–
(rate determining)
(57)
Cl2O3– + ClO2 ¾ Cl2O3 + ClO2–
(58)
Cl2O3 + OH– ¾ HOCl + ClO3–
(59a)
or Cl2O3 + H2O ¾ HOCl + ClO3– + H+
(59b)
ClO2– + H+ ¾ HClO2
131
(fast equilibrium)
(60)
HOCl + HClO2 ¾ Cl2O2 + H2O
(61)
Cl2O2 + H2O ¾ ClO3– + Cl– + 2H+
(62)
This reaction mechanism takes into account the fact that ClO2 reacts faster with hypochlorite ion
than with hypochlorous acid. Thus the reaction takes place at an increased rate at higher pH values
(see the distribution graph of FAC in Chapter 1).
During this reaction several intermediates (Cl2O3–, Cl2O3, and Cl2O2) are formed which are very
reactive, transient species. These intermediates are strong oxidants. The presence of such species may
result in a significant improvement in the disinfectant capabilities of FAC–ClO2 mixtures as compared
with the disinfectant properties of the individual disinfectants.
4.1.2. The C×T principle
The C×T factor is widely used for determining and comparing disinfectant efficacy4. It is based
on the Chick-Watson law122, 123. Here, C is the residual concentration of the disinfectant in mg/L and
T is the time in minutes during which the disinfectant is in contact with the water which is disinfected.
This theory assumes that the reaction between the disinfectant and microorganism is first order with
respect to the disinfectant concentration. Based on this theory the disinfection time is linearly
dependent on the concentration of the disinfectant and this is the generally observed trend for simple
disinfectants. However, when mixed disinfectants are used, deviation from this theory can be
observed and generally suggests interaction between the disinfectants.
The C×T values are practical parameters which need to be determined for each pair of
microorganism and disinfectant, no theoretical method is known for their a priori calculation. These
values are usually given for a certain percent of bacteria killed or more often in log kill units. For
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example, 1 log kill means that 90% of the original bacteria have been killed. The efficacy of the
various disinfectants may change within wide limits for various microorganisms. For example
dissolved chlorine is effective against viruses, but has less efficacy against Giardia lamblia cysts4. The
required C×T value of dissolved chlorine to achieve a 2-log inactivation of viruses is 3 mg/L×minute.
The C×T value for the same inactivation for Giardia cysts is 69 mg/L×minute.
The increased disinfectant efficacy of the mixed disinfectants may occur due to two main reasons.
One reason is synergism between the disinfectants. A possible explanation for synergism is the
following. One of the disinfectants weakens, but does not kill the microorganism and the second
disinfectant can effectively kill this weakened microorganism. The other possibility for the increased
efficacy is the formation of reactive intermediates in a reaction between the two disinfectants. These
intermediates can act as an additional disinfectant. The most important difference between the two
cases is that the intermediates increase the efficacy of the mixed disinfectant only if the two
disinfectants are mixed and used simultaneously. On the other hand, synergism is possible even if the
two disinfectants are applied successively. This latter example is generally used in water treatment,
where a strong disinfectant (ozone or ClO2) is used first and followed by a disinfectant which provides
a residual (dissolved chlorine or chloramines)4, 124-126. Several papers124-126 demonstrated the increased
efficacy of the disinfectants in this type of process.
4.1.3. Microbiological tests
Every disinfectant must pass a series of microbiological tests to be approved by the FDA as a
disinfectant. The details of these tests are defined127 by the AOAC. These tests are very expensive
because they take a long time to perform (incubation period of the microorganism to check if growth
is observed or not) and are very labor intensive. Thus, if any method can be devised to lower the
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number of required tests in the development phase of the disinfectant solutions reliably, the final costs
of the development of the disinfectant solution can be lowered significantly.
The first of these AOAC microbiological tests is a set of laboratory in vitro tests. In the full test,
four different types of bacteria are used on various carriers. The efficacy of the disinfectant is tested
by using 60 of the spore labeled carriers and three different lots of the disinfectant. However, in the
development phase of the disinfectant solution, the cost of testing can be reduced by using only one
lot of the disinfectant and testing only with the most resistant bacteria128.
As the most resistant bacteria, Clostridium sporogenes spores were used on sterile penycilinder
carriers in the development phase. Penicylinders were labeled by soaking them for one minute in a
soil extract nutrient broth containing the spores. The labeled cylinders were dried in vacuum and at
ambient temperature in a desiccator for 24 hours.
Ten mL of the disinfectant solution is added to a 38×200 mm test tube. Five carriers are placed
carefully in the disinfectant, making sure that they do not to touch the sides of the test tube. After a
given time the penicylinders are removed from the disinfectant solution and the kill rate is determined.
The kill rate in the original AOAC test is determined by the number of sterile penicylinders. After
the disinfection process, the penicylinders are rinsed with water and added to a recovery medium. The
spores are removed from the penicylinders by a vortex mixer into the recovery medium. The recovery
medium is incubated for 24 hours. If no bacterial growth is observed after this time, the penicylinder
is marked as sterile. This method for the determination of kill rate gives a straightforward yes/no
answer. This method does not allow accurate determination of the efficacy of the disinfectant because
any cylinder with at least one viable spore would be registered as non-sterile128.
134
However, during the development phase, counting the number of killed spores on each
penicylinder is also possible. Of course, this number gives only the magnitude of the surviving spores.
This number is determined by making serial ten-fold dilutions of the recovery medium. The diluted
solutions are incubated and the number of colonies of the bacteria is counted. The number of colonies
is multiplied by the appropriate dilution factor to get the number of surviving bacteria. This method
is more advantageous in the development phase because it provides more accurate values of the kill
rate than the previous method.
4.2. Experimental
This section gives the details of procedures and chemicals which are first used for this work. All
other procedures are given in the Experimental sections of the previous Chapters.
4.2.1. Preparation of the disinfectant (ClO 2, FAC) solutions
This section describes the generation method of the disinfectants for the initial experiments. The
details of generating these solutions at testing laboratories (microbiological laboratory, corrosion
testing) are given later in the text. The development of new generation methods was necessary,
because microbiologists in general, may not have enough chemical experience in handling ClO2
solutions. Considerable experience is required to be able to use more advanced generation methods
safely. The generation of both ClO2 and chlorine solutions has been described in Chapter 2.
4.2.2. Analytical methods for the measurement of the various species
Two methods can be practically used for the measurement of the various chlorine species in this
system. Chlorine dioxide, hypochlorous acid, hypochlorite ion, and chlorite ion have a characteristic
135
absorbance in the UV-visible range, which can be used for measuring these species simultaneously.
The other method is the iodometric titration.
The application of these two analytical methods is complimentary because of their different
properties. These properties include the speed or the working range of the methods. The UV-VIS
spectrophotometric method has the advantage of being fast and simultaneously providing the
concentration of each species. However, the presence of organic molecules which absorb in the UV
range, may interfere with the measurement of both FAC species and chlorite ion because these species
have their maximum absorbance in the UV region. Furthermore, the useful concentration range is
limited by the range of absorbance values (usually between 0.1 and 2.0 absorbance units) which can
be measured accurately. Spectrophotometric measurements are well suited for kinetic measurements.
In contrast, the iodometric titration is slow. It may take up to 30 minutes to determine the
concentration of all three species, mainly due to the purging step (to remove ClO2) and the reaction
time required for the formation of iodine. This method is a differential method which means that the
errors are cumulative (see Chapter 2 for details). The concentration range of iodometric titrations can
be relatively wide. By changing the sample size, the maximum or minimum analyte concentration can
be adjusted as required. The detailed description of these methods is given in the following
paragraphs.
4.2.3. Iodometric titration
Iodometric titrations were performed based on previously described procedure9, 30, 43 which
utilizes the pH dependent reactivity of the various chlorine species. The chemical background for the
titrations is given in Chapter 2. To determine the three species present (FAC, chlorite ion, and ClO2),
136
four steps were necessary. This method is not able to distinguish the two FAC species (HOCl and
OCl–) from each other.
1.
This step measures the concentration of dissolved chlorine and 1/5 of the ClO2 concentration.
The sample was added to a pH 7.0 phosphate buffer and potassium iodide was added. The
iodine formed was titrated to the end point by using standard sodium thiosulfate solution.
2.
The same sample was used for the second step. This step measures chlorite ion and 4/5 of
ClO2. To the sample 2.5 M HCl solution was added and the solution was allowed to react in
the dark for five minutes. It was again titrated to end point with standard sodium thiosulfate
solution.
3.
For this step, a new sample was used which was added to a pH 7.0 phosphate buffer in a
beaker. The solution was purged for fifteen minutes with pre-purified nitrogen gas. This
purging removes ClO2 and partially dissolved chlorine. Potassium iodide was added to the
solution and the iodine formed titrated to the end point. The volume of Na2S2O3 in this
titration was not used for further calculations, but it was necessary to remove any remaining
dissolved chlorine from the sample.
4.
In this step, the same sample was used as for step three. The sample was acidified with 2.5 M
HCl and allowed to react for five minutes. The iodine formed was titrated with standard
sodium thiosulfate solution.
These titrations can be performed by using an automatic titrator, such as a Radiometer
autotitrator. The titrator is able to perform complex titration sequences, in which more than one
reagent is used, wait times (to allow some reactions to be completed) are included, etc. The results
of such titrations are not displayed on the screen of the titrator, but they are sent to a computer
through a standard RS232 interface. Writing a short program in MS QuickBASIC was necessary to
capture these results on the computer. The source code of this program is given in Appendix A. The
program captures all titration data (including all titration points and end points) and saves them in a
text file.
137
4.2.4. Spectrophotometric measurement
Spectrophotometric measurements were used to follow the concentration changes during the
kinetic measurements. This method can be used to determine rapidly the concentrations. Rapid
determination is necessary in the kinetic measurements due to the fast reaction. However, this method
is limited to measuring relatively low concentrations due to the limitations in the absorbance
measurements. For all spectrophotometric measurements, a tandem cell was used (see Chapter 3 for
details).
To use the spectrophotometric measurement, the molar absorptivities of the different species had
to be determined. The method for the determination of molar absorptivities for chlorite ion and ClO2
is described in Chapter 3. The details for determining the molar absorptivities of HOCl and OCl– are
given here8.
For reliable molar absorptivity measurements, the pH of the FAC solution has to be chosen such
that only one species is present predominantly (>99%). If this is not true, the exact value of the
equilibrium constants and pH is necessary for accurate measurements. This would make this
measurement prone to significant error. The molar absorptivity of HOCl was measured by adjusting
the pH to 3. The molar absorptivity of OCl– was measured at pH 10. The concentration of the FAC
was varied between 4.0×10–4 M to 4.9×10–3 M in both cases. The FAC solutions were titrated prior
to the spectrophotometric measurements.
The molar absorptivities of the various species are shown in Figure 39. Table 26 compares the
molar absorptivities of the various chlorine species at the wavelengths of their maximum absorbances.
From these data and the figure, it can be concluded that the peaks of the species are well separated.
138
Thus, determining the concentration of each of these species by spectrophotometric measurement is
possible in the presence of the other species.
Figure 39. Comparison of the molar absorptivities of the
various chlorine containing species in the mixed
disinfectant system. — Hypochlorous acid, — chlorite
ion, — hypochlorite ion, — ClO2
Table 26. Comparison of the molar absorptivities of the various chlorine
containing species at the wavelengths of the maximum absorptivities.
Bold numbers indicate the maximum molar absorptivity (M–1cm–1) for
the given species.
235 nm
260 nm
292 nm
358 nm
Hypochlorous acid
90.6
37.7
23.5
0
Chlorite ion
69.2
148.2
90.6
2
Hypochlorite ion
9.6
102.8
343.6
12.3
167.4
55.9
180.5
1128.9
ClO2
The software provided with the spectrophotometer has a function which can be used to
determine the concentration of various species in a mixture based on the spectra of standard solutions.
In the initial step, the spectra of the standard solutions are loaded and the concentration of the various
139
species in each spectrum is entered. The calibration can be performed either for a wavelength range
or a set of separate wavelengths (for example the wavelengths of the peaks of the various species).
Based on these data, the program creates a calibration which can be used for calculating the
concentrations of the different species in an unknown mixture. The advantage of this method is that
it can use a large number of wavelengths for determining the concentrations, possibly resulting in
more accurate values. However, if the spectra of the species partially overlap in a complex mixture
(such as the current system), the software may determine inaccurate (e.g., negative) concentrations.
Furthermore, transferring the determined concentrations to other programs for further calculations
can be problematic.
The concentration of the various species can be determined by using Excel. In this case only the
wavelengths of the maximum absorbances are used and the following linear equation system is solved
by using Gauss elimination107. A short program has been written for this purpose in VBA.
(63 a)
(63 b)
(63 c)
(63 d)
4.3. Preparation of the Mixed Disinfectant Solutions
According to DOT regulations22 ClO2 can not be shipped, it needs to be generated at the point
of application. The normal methods used to generate ClO2 require significant chemical expertise and
experience with ClO2. Thus it is not suitable for generating ClO2 at a laboratory where the researchers
140
(microbiologists) have limited chemical experience. The microbiological tests, which are required for
the approval of a disinfectant, are typically performed by an independent microbiological laboratory.
This laboratory needs to have access to the mixed disinfectant solutions to perform the required tests.
The disinfectant solutions have a short life-time due to the reaction between FAC and ClO2. Thus,
they can not be shipped in mixed form and they need to be prepared before testing. Therefore
devising a simple method for the generation of ClO2 was necessary. This method can be performed
by “almost anybody” with some chemical experience.
The generation method should work in a consistent and reliable way, providing disinfectant
solutions with constant composition. In this way, the generated disinfectant solutions would always
have the same and known concentration of ClO2 and FAC. A constant, but unknown concentration
of other species would be ensured this way. These other species could be for example chloride and
carbonate ions. These ions can have a significant influence on the rate of the FAC-ClO2 reaction18 and
on the disinfection efficacy of the mixed disinfectants. Therefore, any change in the composition of
the mixed disinfectant solutions can lead to unexpected and hard to explain results.
When the effect of various parameters of the mixed disinfectant is determined and the
microbiological efficacy is known, it would be possible to generate the mixed disinfectant solutions
by an electrochemical generator. This generation method has many advantages. One important
advantage is that both ClO2 and FAC can be generated simply from chemicals which require minimal
safety precautions. Chlorine dioxide and dissolved chlorine can be generated effectively from sodium
chlorite and sodium chloride, respectively. In this case only solid chemicals need to be transported
to the application site. Sodium chloride can be transported in large amounts without complications.
Sodium chlorite is more problematic, but if the necessary precautions are taken, sodium chlorite can
141
by shipped safely. Furthermore, by using electrochemical generators, the amount of disinfectant
generated can be controlled conveniently to meet the demand.
However, the purpose of the current research was not to use this generation method, which
ultimately would be a commercial process. The reason is that this research was intended to explore
the interactions in the disinfectant system and how the different parameters (disinfectant
concentration, pH, and temperature) affect the disinfection capabilities of the mixed disinfectant
solutions. Thus, ClO2 and FAC were generated by chemical methods. These chemical methods
provided the required disinfectants in a simple way and with a consistent composition. If
electrochemical generators had been used, determining how the various generation parameters
affected the composition of the generated disinfectants would have been necessary. Determining the
relationship between the generation parameters and the composition of the disinfectant solution
would require significant time and effort. However, at the initial phase, this time and effort did not
appear to be justified, because at that point it was not proven that the mixed disinfectant solutions
would provide the expected efficacy.
The methods, which are described here, can be easily used for the generation of the disinfectant
solutions in a laboratory. However, typically they are not practical for commercial applications. For
practical applications of the mixed disinfectant solutions, finding alternative methods for the
generation of these disinfectants is necessary. These methods would provide the necessary solutions
efficiently in large volume.
Chlorine dioxide can be effectively generated from sodium chlorite through chemical
reactions2, 10-12, 23, 31. One such method has been described in Chapter 2 (the reaction of NaClO2 and
K2S2O8). However, that method requires some chemical expertise and does not meet the previously
142
described requirements. Thus, in this application, two main methods were considered: 1) reacting
chlorite ion with FAC; 2) reacting chlorite ion with a strong acid. If excess FAC is used, the
advantage of the first method would be that the mixed oxidant solutions are generated in one step.
Whereas the second method would give “pure” ClO2 solutions. These solutions would be buffered,
diluted, and mixed with an FAC solution to prepare the mixed oxidant solutions. In the following
sections these two ClO2 generation methods are discussed.
The generation methods presented are useful only for the generation of ClO2 on a small scale and
not intended as commercial generation methods. Their purpose is to provide the independent
microbiological testing laboratory with sufficient amounts of mixed disinfectant solutions.
4.3.1. Generation of ClO 2 by mixing FAC and chlorite ion
The reaction of chlorite ion with chlorine –either in gaseous or aqueous phase– is very often used
for generating ClO2. This method could be advantageous for preparing mixed disinfectant solutions.
If excess FAC is used to generate ClO2, some of the original FAC would remain. Thus, this method
would provide the mixed disinfectant solutions in one step. To achieve this goal, it is necessary to
generate the ClO2 at a pH value which is used in the disinfectant solutions. The reason for this is the
relatively fast reaction between ClO2 and FAC. If ClO2 is generated at a different pH value, during
the time required to buffer the generated solution, a significant amount of the disinfectants would be
lost.
The chlorine–chlorite ion reaction is generally used at acidic pH values. Thus, it was necessary
study how effectively the reaction generates ClO2 at the currently used pH values of 7.0, 7.5, and 8.0.
This pH values were maintained by phosphate buffers.
The reaction is described with the following general equations9:
143
2 ClO2– + Cl2(g) = 2 ClO2 + 2 Cl–
(4.a)
2 ClO2– + HOCl = 2 ClO2 + Cl– + OH–
(4.b)
The mechanism of these reactions can give very important information about what parameters
favor the formation of ClO2. The detailed mechanism of the reaction is given in Equations 5-7.
Cl2 + ClO2– Ž [Cl2O2] + Cl–
(5)
2 [Cl2O2] Ž 2 ClO2 + Cl2
(6 a)
[Cl2O2] + ClO2– Ž 2 ClO2 + Cl–
(6 b)
[Cl2O2] + H2O Ž ClO3– + Cl– + 2 H+
(7)
The formation of ClO2 is favored when the intermediate, Cl2O2, is formed in high concentration.
If the concentration of the intermediate is low, chlorate ion is the major product of the FAC-chlorite
ion reaction.
Most of the experiments were performed with excess FAC, as it would be needed for the one
step preparation of the mixed disinfectant solutions. Other experiments were performed with excess
chlorite ion to test the effect of the initial reactant concentration ratio.
Based on the results of these measurements, the following can be concluded. The amount of
generated ClO2 is highly dependent on the pH. This is illustrated in Figure 40. At pH 7.0, ClO2 is
generated rapidly and reaches its maximum concentration in a few seconds. The generated ClO2
rapidly disappears from the system by means of the reaction with FAC. As the pH increases the rate
of generation of ClO2 slows and the highest concentration of ClO2 achieved decreases. However, at
these pH values the ClO2–FAC reaction takes place slower, resulting in higher “final” ClO2
concentration.
144
Figure 40. The formation of ClO2 with excess FAC as a
function of time at various pH values. [FAC] = 1.07×10–2 M,
[ClO2–] =5.97×10–3 M, pH =  7.0, Ž 7.5, • 8.0
Figure 41. The formation of ClO2 with excess chlorite ion as
a function of time at various pH values. [FAC] =
1.07×10–2 M, [ ClO2–] = 1.79×10–2 M, pH =  7.0, Ž 7.5, •
8.0
The amount of generated ClO2 and the time required to reach the maximum concentration is
influenced by the chlorite ion to FAC ratio. This can be clearly seen by comparing Figures 40 and 41.
When the FAC is in excess, the time required to reach the maximum ClO2 concentration is
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significantly dependent on the pH. However, when the chlorite ion is in excess, the variation in the
time required to reach the maximum concentration of ClO2 is much smaller.
Figure 42 shows three aspects of the ClO2 generation which need to be considered. First, it
shows that at higher pH values, higher chlorite ion concentration is needed to reach similar ClO2
concentrations (the difference is ~10%). Second, it shows that it takes longer time to reach the
maximum ClO2 concentration at higher pH. From the figure it is clear that ClO2 is reacting with FAC
at a slower rate at higher pH values.
Figure 42. The formation of ClO2 with excess chlorite ion as
a function of time. The chlorite ion concentrations are adjusted
to reach similar ClO2 concentrations. At pH 7.0 (), [FAC] =
1.07×10–2 M, [ClO2–] = 1.79×10–2 M; at pH 7.5 (Ž), [FAC] =
1.06×10–2 M, [ClO2–] = 2.99×10–2 M
The effect of initial chlorite ion concentration on the maximum ClO2 concentration generated
revealed a linear relationship, as shown in Figure 43. The figure shows the concentration of the
maximum ClO2 concentration generated as a function of chlorite ion concentration. This maximum
146
concentration was reached at different time intervals after mixing for different chlorite ion
concentrations. The Figure includes the ClO2 concentrations which were reached after five and ten
minutes.
Figure 43 shows that the concentration of the generated ClO2 is linearly dependent on the
chlorite ion concentration only when the maximum ClO2 concentration generated is considered. The
other two curves deviate markedly from the least squares fitted lines. This non-linearity is due to the
reaction between the generated ClO2 and FAC. At higher chlorite ion concentration the ClO2
concentration generated is higher. However, this higher ClO2 concentration results in a more rapid
disappearance of the ClO2 due to the FAC–ClO2 reaction.
Figure 43. The dependance of the concentration of the generated
ClO2 on the chlorite ion concentration at constant FAC concentration.
pH = 7.0, FAC = 1.07×10–2 M,  Maximum ClO2, Ž after 5 minutes,
• after 10 minutes. The equation of the least squares fit for the
maximum ClO2 concentration: c(ClO2, M)max = 0.14×c(ClO2–, M) –
2.4×10–4
147
This linear relationship of the maximum ClO2 concentration holds at each tested pH value.
Increasing the initial chlorite ion concentration to a level which would generate the ClO2
concentration required, results in lower FAC concentration than required. In order to have the
required FAC concentration in the mixed disinfectant solution, it is possible to use higher initial FAC
concentration or add more FAC at the end of the reaction. In the first case the higher FAC
concentration would result in even higher ClO2 consumption, thus the ClO2 concentration required
may not be reached. The latter case does not provide any advantage over other ClO2 generation
methods, because it would be a two-step process as the other methods.
The generation of ClO2 with the reaction of chlorite ion with FAC is not able to generate the
mixed disinfectant solutions in a single step at the required concentrations. Thus, this method does
not have any advantage over other chlorite ion based ClO2 generation methods.
4.3.2. Generation of ClO 2 by mixing sodium chlorite with strong acid
Chlorous acid is a relatively strong acid (pKa~1.7) which undergoes a fast selfdecomposition24, 25. The products of the decomposition are ClO2, chloride, and chlorate ions. The
distribution of these products is dependent on the parameters of the reaction (pH, chlorite ion and
chloride ion concentration). Under appropriate conditions this method can provide high enough ClO2
concentrations. The reaction is described in a simple form as
4 HClO2 Ž 2 ClO2 + ClO3– + Cl– + 2 H+ + H2O
(9)
5 HClO2 Ž 4 ClO2 + Cl– + H+ + 2 H2O
(10)
For acidifying the chlorite ion solution strong acids are necessary, because chlorous acid itself
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is a relatively strong acid. Sodium bisulfate and hydrochloric acid were tested for generating ClO2 in
this research, because they are readily available.
4.3.3. The reaction of chlorite ion with sodium bisulfate
The concentration of bisulfate ion tested ranged from 0.1 M to 0.8 M. Chlorite ion
concentrations varied from 0.003 M to 0.24 M.
When low concentrations of chlorite ion (below 0.03 M) were used, less than 50 mg/L ClO2 was
generated. Upon increasing the chlorite ion concentration, the concentration of the ClO2 formed is
increased. But it did not reach sufficiently high concentrations even at the highest chlorite ion
concentration. Furthermore, the reaction time to generate the maximum ClO2 concentration is long,
about 15-20 minutes.
To reach high enough ClO2 concentration, using more concentrated chlorite ion solutions would
be necessary which could present a severe safety problem for inexperienced users. Therefore, this
ClO2 generation method was not studied further.
4.3.4. The reaction of chlorite ion with hydrochloric acid
The concentration of HCl used ranged from 0.01 M to 0.8 M in these experiments. The chlorite
ion concentrations were between 0.003 M and 0.24 M. The reaction was initially followed by
iodometric titration and later by spectrophotometric measurements.
Low concentrations of chlorite ion with low hydrochloric acid concentrations did not generate
detectable amounts of ClO2 within 15-20 minutes. Increasing the acid concentration resulted in fast
ClO2 generation which was complete in about five minutes. In this case the concentration of the
generated ClO2 was linearly dependent on the chlorite ion concentration (see Figure 44).
149
The concentration of chlorite ion solution, which is required for generating concentrated ClO2
stock solutions, can be determined based on the linear relationship. These ClO2 stock solutions can
be buffered and diluted, mixed with buffered FAC solutions to prepare the required mixed disinfectant
solutions. The ClO2 stock solution generated this way is strongly acidic. Thus, a high buffer capacity
is necessary to adjust the pH of the final solution to the selected pH values. Therefore, a fairly
concentrated ClO2 stock solution of about 6 g/L is required. When this solution is diluted to reach
the ClO2 concentration required, the acidity is also decreased. In this way, the required buffer capacity
can be lowered. About 0.25 M chlorite ion is required to reach the 6 g/L ClO2 concentration. This
solution is relatively dilute, thus it presents only minimal safety problem. The chlorite ion solution is
mixed with an equal volume of 1.0 M hydrochloric acid solution. Five minute reaction time is required
to generate the ClO2 stock solution.
Figure 44. The dependence of the ClO2 concentration
generated on the initial chlorite ion concentration.
c(HCl) = 0.5 M The equation of the linear fit: c(ClO2, mol/L)
= 0.36×c(NaClO2–, mol/L) + 2.7×10–4
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4.3.5. Preparation of FAC solutions
The selection of a method for the generation of FAC solutions for the laboratory preliminary tests
proved to be simple, because the commercially available bleach solutions contain sufficiently high
dissolved chlorine concentrations. The concentrations of these solutions were determined by
iodometric titrations and were diluted. The pH of the diluted solutions was adjusted to 11 to reduce
the decomposition8 of FAC.
However, these bleach solutions contain other, potentially interfering chemical species (e.g.,
chloride ion and surfactants). To ensure that these species do not significantly affect the kinetic
parameters of the FAC–ClO2 reaction, subsequent kinetic runs were performed by using these FAC
solutions.
4.3.6. Preparation of phosphate buffers
The buffers generally used have low buffer capacity, which is not enough to adjust the pH of the
disinfectant solutions to the required value. For this reason special, concentrated phosphate buffers
were needed to prepare the mixed oxidant solutions. Concentrated sodium dihydrogen phosphate
solutions (0.5 M) were “titrated” with 1.0 M carbonate free sodium hydroxide solutions until the
required pH was reached.
4.3.7. Preparation of mixed disinfectant solutions
Once the individual disinfectants can be generated with consistent composition, a method is
needed to mix them in order to obtain the mixed disinfectant solution. In this research, the approach
was to generate concentrated stock solutions of the individual disinfectants, buffer and dilute them
151
to the required pH and concentration and mix them to prepare the mixed disinfectant solutions. To
achieve this, the following method was devised.
Concentrated ClO2 solutions are prepared by mixing equal volumes of 1.0 M HCl solution and
0.25 M NaClO2 solution. The solutions are allowed to react for five minutes, after which time the
solution contains ~6 g/L ClO2. Aliquots of this solution are added to phosphate buffers in volumetric
flasks and diluted to volume.
Concentrated FAC stock solutions are prepared by diluting the commercially available bleach
solution and adjusting its pH to 11. Aliquots of this FAC stock solution are added to phosphate
buffers in volumetric flasks and diluted to volume. Both FAC and ClO2 solutions are used
immediately after preparation. Mixed disinfectant solutions are prepared by mixing the diluted FAC
and ClO2 solutions in equal volumes.
4.4. Initial experiments
The initial experiments were necessary to determine the suitable pH and concentration of the
disinfectants to optimize the rate of the reaction between FAC and ClO2. These experiments were
used to determine whether the reaction takes place on a time-scale which would make the
microbiological experiments feasible.
The pH of the disinfectant is important for two reasons. It affects the rate and mechanism of the
reaction between FAC and ClO2 (Eq. 55) and in addition, the pH influences the efficacy of dissolved
chlorine as a disinfectant.
2 ClO2 + HOCl + H2O = 2 ClO3– + Cl– + 3 H+
152
(55)
With increasing pH, the rate of the reaction increases and the efficacy of FAC decreases. The
increased rate has a negative effect on the disinfection efficiency of the mixed disinfectant solution,
because the initial disinfectants are used up rapidly. However, if the reaction takes place slowly, the
concentration of the intermediates is always low during the reaction, thus their contribution may be
negligible in the disinfection. Therefore, selecting carefully the pH of the disinfectant solution is
necessary to balance these factors.
The pH of the solutions initially tested were 9.0 and 7.0. The concentration of both FAC and
ClO2 in these solutions was 200 mg/L. In these experiments the reaction was followed by iodometric
titrations. At pH 9.0 the reaction is too fast to be used for the mixed disinfectant solutions. At pH 7.0
the half-life of the reactions was about 15 minutes. Thus, by using 200 mg/L of both disinfectants,
the rate of the reaction is in the range that would be considered useful.
Based on these results, pH values around 7.0 appear to meet the previously described criteria for
a moderately fast reaction between FAC and ClO2. The pH values for further testing were determined
on the following basis. At pH 7.5 the ratio of HOCl and OCl– is 1:1, because the pKa of hypochlorous
acid is about 7.5. Close to this pH value, small changes in the pH result in significant changes in the
ratio of HOCl and OCl–. This difference in the relative concentration of the two FAC species would
determine whether the contribution of these species to the observed rate of the FAC–ClO2 reaction
or to the disinfection process is different. If there were only a small change in the relative
concentration of HOCl and OCl– at the two selected pH values, the variation in the observed rate at
the two pH values would be small — even if the contribution of the two species to the FAC–ClO2
reaction or the disinfection process is different. However, knowing these contributions is important
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in order to be able to improve the efficacy of the disinfectant. Thus, the two selected pH values were
6.5 and 7.5.
4.5. Kinetic study
The kinetic studies were performed to determine the rate equation and the kinetic parameters
(i.e., rate constant, activation enthalpy, entropy). These studies were necessary because the mixed
disinfectant solution is a complex mixture, which contains many other species besides the reactants.
Many of these species may have significant influence on the mechanism and the rate of the FAC-ClO2
reaction. The purpose of these measurements was to confirm that the FAC-ClO2 reaction takes place
according to the previously published studies17, 18. Furthermore, these kinetic results can be used in
interpreting the initial microbiological results and even be used in a predictive manner in order to
eliminate some of the subsequent microbiological tests. This would result in lower cost of the
development of the new disinfectant.
Based on the results of the initial studies, the reaction was followed at pH 6.5 and 7.5. The
concentration of ClO2 was varied from 12.5 mg/L (1.85×10–4 M) to 200 mg/L (2.97×10–3 M). The
FAC concentration was varied from 50 mg/L (7.04×10–4 M) to 300 mg/L (4.23×10–4 M). These
concentrations correspond to FAC to ClO2 ratios from 1:2 to 24:1. The measurements were
performed at 22°C and 35°C. The reaction was followed spectrophotometrically in a tandem cell (see
Chapter 3 for details on the cell). One compartment contained FAC and the other compartment ClO2
solution.
154
The concentration of chlorite ion (as a product), FAC, and ClO2 are shown for a typical run in
Figure 45. The concentration of chlorite ion goes through a maximum which is due to the fact that
chlorite ion reacts with FAC to form ClO2.
Figure 45. Concentration changes of the main species in the
FAC–ClO2 system. [ClO2] = 1.48×10–3 M (100 mg/L), [FAC] =
2.12×10–3 M (150 mg/L), pH = 7.5, temperature = 22 °C. 
Chlorite ion, Ž ClO2, • FAC
To determine the rate equation, the method of initial rates was used129. The details of this method
are given below. The generalized form of the rate equation is
(64)
The initial rate is the rate of the reaction at the beginning of the reaction. In practice the initial
rate is measured when the extent of the reaction is <10%. In this case the concentration of the
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reactants is not significantly different from their initial concentrations. The initial rate was determined
by fitting the ClO2 and FAC concentration changes with polynomial curves.
If the concentration of one of the reactants (e.g., FAC) is held constant, the initial rate (ri) is,
(65)
where
If the concentration of the ClO2 is varied, Equation 65 can be linearized by taking the logarithm
of this equation. The slope of the resulting line gives the order with respect to ClO2 (x).
(66)
The order of the other reactant (FAC) can be determined by repeating this procedure while keeping
the ClO2 concentration constant.
Figures 46-49 show the results of the fitting. Tables 27 and 28 show the concentration of the
disinfectant, which was held constant, the pH, and the order of the disinfectant in question (the slope
of the linear fit). The determined reaction orders with respect to ClO2 and FAC is 1.1 ± 0.18 and 1.0
± 0.28, respectively. Thus, the rate equation is
(67)
According to the previously discussed reaction mechanism, the two FAC species (hypochlorous
acid and hypochlorite ion) are expected to react at different rates with ClO2. The reaction mechanism
implies that hypochlorite ion reacts faster with ClO2. This expectation is supported by the observed
pH dependence of the reaction rate: the reaction of FAC and ClO2 takes place at a faster rate with
156
Figure 46. Determination of the reaction order of FAC at pH 6.5 by using the method
of initial rates.
Figure 47. Determination of the reaction order of FAC at pH 7.5 by using the method
of initial rates.
157
Table 27. Comparison of the order of FAC at various constant
ClO2 concentrations and pH values.
Constant ClO2
concentration (mg/L)
pH
Reaction order of
FAC
100
6.5
0.81
100
7.5
0.93
50
6.5
1.16
50
7.5
0.7
25
6.5
1.26
25
7.5
1.53
12.5
6.5
0.83
12.5
7.5
0.85
Average
1.0
Standard Deviation
0.28
Figure 48. Determination of the reaction order of ClO2 at pH 6.5 by using the method
of initial rates.
158
Figure 49. Determination of the reaction order of ClO2 at pH 7.5 by using the method
of initial rates.
Table 28. Comparison of the order of ClO2 at various
constant FAC concentrations and pH values.
Constant FAC
concentration
(mg/L)
pH
Reaction order
of ClO2
300
6.5
1.04
300
7.5
0.99
200
6.5
0.86
200
7.5
1.26
150
6.6
0.93
150
7.5
1.12
100
6.5
0.9
100
7.5
1.39
50
6.5
1.16
Average
1.1
Standard Deviation
0.18
159
increasing pH, where the FAC exists predominantly in the form of hypochlorite ion. The rate equation
contains the total FAC concentration. The total FAC concentration is the sum of the concentration
of HOCl and OCl–. Thus, by taking into account the dissociation of hypochlorous acid (Equation 56),
kobs can be expanded the following way:
H+ + OCl– ¾ HOCl
(56)
(68)
(69)
where Kp
kHOCl, kOCl–
the protonation constant of OCl– (log Kp . 7.40)
rate constants for the reaction of ClO2 with HOCl and OCl–, respectively
Table 29 compares these rate constants for the two FAC species at two temperatures. The rate
constant for the ClO2–hypochlorite ion pathway is about three orders of magnitude higher than for
the ClO2 –hypochlorous acid reaction. This means that at pH values at which OCl– is present, the
reaction proceeds almost exclusively through this pathway. Thus, the rate of the overall reaction
primarily depends on the fraction of the FAC which is present as hypochlorite ion.
Table 29. Comparison of the rate constants for the different
reaction pathways and temperatures.
22 °C
35 °C
kOCl–
0.52 ± 0.15 s–1
1.5 ± 0.3 s–1
kHOCl
(1.6 ± 0.1) ×10–3 s–1
(2.1 ± 0.3) ×10–3 s–1
Figure 50 shows the quality of the fit by using this rate equation and rate constants. The Figure
clearly shows that the fit between the measured and calculated concentrations is good.
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Figure 50. Comparison of the measured and fitted ClO2
and FAC concentrations.  Measured FAC concentration,
Ž Measured ClO2 concentration, the solid lines represent
the fitted concentrations
4.5.1. Temperature effect
Determination of the effect of temperature on the reaction is important. If an electrochemical
method is used for generating the disinfectants, the current which passes through the solution would
heat the disinfectant solution. Knowing the temperature effect would allow the correction for the
changes in efficacy. For example, lowering the temperature would slow the reaction, prolonging the
life-time of the mixed disinfectant solution.
Table 29 shows the values of the rate constants at two temperatures. By comparison, at 35°C
the reaction is about three times faster than at 25°C. This means that in a real world disinfection
process, cooling the disinfectant solution may be necessary to keep the rate of the reaction under
control such that high disinfection is achieved.
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The rate constants can be used for determining the activation enthalpy (DH‡) and entropy (DS‡)
by using the Eyring equation:
(70 a)
Alternatively this equation can be plotted in the following form:
(70 b)
To improve the quality of the least squares fitting of Equation 70 b, all calculated rate constants
were used instead of using only the average rate constants. The determined activation parameters are
DH‡ = 64 ± 4 kJ/mol and DS‡ = –34 J/(mol K). These values are in good agreement with the
previously published values, DH‡ = 66.5 kJ/mol and DS‡ = –22.3 J/(mol K). The difference between
the previously determined and the current values are possibly due to the different conditions which
were used in the two works. In the previous work18, pure hypochlorous acid solutions were used. For
this research, the hypochlorous acid solution was prepared from a bleach solution, which contained
other “contaminants,” such as chloride and possibly chlorate ions.
4.6. Microbiological results
The initial microbiological testing was performed at two pH values, 7.0 and 7.5 and with three
different disinfectant solutions. The compositions of these solutions were 100 mg/L FAC and
100 mg/L ClO2, 200 mg/L FAC and 100 mg/L ClO2, and 200 mg/L FAC and 200 mg/L ClO2.
The reason for selecting these three solutions was that in this way the effect of both FAC and
ClO2 concentration on the disinfection efficiency can be determined with a small number of test
162
solutions. Furthermore, the reason for selecting these pH values was that the pKa of hypochlorous
acid is about 7.5, meaning that both hypochlorous acid and hypochlorite ion are present in significant
amounts at the selected pH values. Thus these two values, where the relative abundance of OCl– and
HOCl is different, can give important information about how the changes in the concentration of these
two species affects the efficacy of the mixed disinfectant.
Based on the C×T concept, the disinfection time was expected to be inversely dependent on the
concentration of both disinfectants. Furthermore, the form of dependence of the disinfection time was
expected to be the same at both pH values. However, because of different disinfectant efficacies of
HOCl and OCl–, the disinfection times were expected to differ significantly at the two pH values.
In contrast with these expectations, the microbiological results revealed a more complex
concentration dependence for the disinfection times. In addition, the form of this concentration
dependence changed at the two tested pH values. This complex concentration and pH dependence
indicates the presence of some alternative disinfectant (probably a short-lived intermediate), which
plays an important part in the disinfection process. The fact that the concentration dependence of the
disinfection changes with pH, shows that the formation and/or the disappearance of the alternative
disinfectant species are a pH dependent reaction. The dependance of the disinfection times (tdisinfection)
on the disinfectant concentrations is given below.
at pH 7.0
tdisinfection % [FAC]1, [ClO2]0
(71)
at pH 7.5
tdisinfection % [FAC]2, [ClO2]2
(72)
Based on the results of the microbiological tests, the following conclusions can be drawn. At pH
7.0, the disinfection time is inversely proportional to the concentration of FAC as is expected from
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the C×T model. However, in contrast with the theory, the disinfection time is independent of the
concentration of ClO2 in the tested concentration range.
As it has been described earlier, the OCl––ClO2 reaction is faster than the HOCl–ClO2 reaction
and the formation of the intermediate(s) shows the same rate dependence. At pH 7.0, however, the
concentration of OCl– is low. Thus the concentrations of the intermediates are mainly controlled by
the concentration of hypochlorite ion and practically independent of ClO2 concentration which is in
significant excess. As the pH 7.5 case suggests, the intermediates have an important role in the
disinfection process. Thus if the formation of the intermediates is only slightly influenced by the ClO2
concentration, the overall disinfection can be independent of the ClO2 concentration in a small range.
At pH 7.5, the disinfection is dependent on the concentration of both disinfectants as is expected
from the underlying chemistry. However, in contrast to these expectations, the disinfection time is
dependent on the square of the concentrations of both disinfectants. This form of concentration
dependence suggests that alternative disinfectants (i.e., intermediates) are contributing significantly
to the disinfection process.
During the initial testing, no control experiments were used. However, it was known from earlier
microbiological experiments that the disinfection time of only FAC at 650 mg/L is more than 24
hours. In the current test, the disinfection was achieved within 1-2 hours. This clearly shows that the
mixed disinfectant solutions greatly outperform the simple chlorine-based disinfectant solution, even
when the lowest of the proposed concentrations of FAC and ClO2 are used. Thus, the mixed
disinfectant solutions have significant advantages over the traditional, chlorine-based disinfectants.
Very fast disinfection is achieved by using low concentration of disinfectants, which means that the
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disinfectant solutions present less safety problem and are less corrosive due to the lower disinfectant
concentrations.
To improve the efficacy of the mixed disinfectant solutions, the concentrations of both
disinfectants were increased and further microbiological studies were performed. The composition
of these new disinfectants is summarized in Table 30.
Table 30. The composition of mixed disinfectant
solutions for the second microbiological studies.
#
FAC
concentration
(mg/L)
ClO2
concentration
(mg/L)
1
400
200
2
200
400
3
300
300
4
400
400
Even though the first two solutions are expected to have the same disinfection time, they may
not be equally suitable disinfectant solutions. The reason for this can be understood from the
following facts. FAC solutions are known to be corrosive, on the other hand ClO2 solutions are
considerably less corrosive. Thus, increasing the concentration of FAC in the mixed disinfectant
solution increases the corrosive properties of the mixed disinfectant solutions. In contrast, ClO2 is
more volatile than FAC. Therefore, increasing the ClO2 concentration in the disinfectant may result
in a significant increase in the health hazard of the disinfectant solution for the users. For these
reasons it is necessary to consider carefully these options and find a disinfectant mixture which
satisfies a preset list of requirements.
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4.6.1. Second set of microbiological studies
The disinfectant solutions, summarized in Table 30, are expected to have disinfection times on
the order of a few minutes. These estimates are based on the findings of the first set of
microbiological experiments.
The results of the new microbiological studies, however, did not show the expected efficacy.
According to these results, hardly any disinfection took place in the first 30 minutes. Even after this
time, the disinfection efficacy of the concentrated mixtures was not superior to the efficacy of the
mixed disinfectants at lower disinfectant concentrations. This deviation from the previously observed
disinfection model indicated that at high disinfectant concentrations other effects were contributing
to the disinfection efficacy measured by the AOAC testing procedure.
The results showed low disinfection in the first 30 minutes of the disinfection. This is in contrast
with the predictions of the chemical kinetics. The concentration of the various intermediates is the
highest during this initial period. Thus, significant disinfection efficacy would be expected during this
time. However, the results showed that for some reason the high concentration of the disinfectant and
intermediates did not result in effective disinfection.
From the microbiological experimental procedure, it is well-known that the tested penicylinders
are covered with a high amount of protein. Furthermore, discussion with the supervisor of the
microbiological testing laboratory128 revealed that this protein coating requires about a 30 minute
wetting time. During this wetting time, the efficacy of the disinfectants can be hindered because the
interaction of the spores and the disinfectant is minimal.
When high disinfectant concentrations are used, more than 95% of the ClO2 is used up in the first
30 minutes. This means that after this time, the only disinfectant left in the solution is FAC, therefore
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the mixed disinfectant shows only the efficacy of FAC. Thus, the FAC–ClO2 reaction may consume
the disinfectants during the time when only minimal disinfection can take place, resulting in no
improvement in the efficacy of the mixed disinfectants. In addition, the proteins on the cylinders can
show demand for both ClO2 and FAC. At high ClO2 concentrations, evaporation loss is possible. This
is especially important in the AOAC test127, where the used test tubes have large headspace.
4.7. The effect of the penicylinders on the FAC–ClO2 reaction
Further kinetic studies were performed to determine the effect of penicylinders on the FAC–ClO2
reaction. The studies were conducted by using sterilized penicylinders. These cylinders were prepared
according to the procedure described127 by the AOAC, but the spores were killed by heating the
cylinders to 121°C. The disinfectant solutions were added to vials which contained the penicylinders.
These vials had sufficiently large headspace to have the same conditions as in the AOAC tests.
Initially the concentration of the disinfectants was followed by spectrophotometric
measurements. “Blank” experiments were performed where the mixed disinfectant solutions were
added to vials without a penicylinder.
During the initial work, an absorbance increase was observed in spectra of the solutions in the
200-260 nm region. This absorbance increase interferes with the spectrophotometric measurement
of chlorite ion and FAC species. Thus in these studies only ClO2 concentrations were followed. The
absorbance increase is possibly due to the release or solubilization of some of the proteins from the
penicylinders. This may be the most reasonable explanation because the absorbance increase was
observed in the case when the penicylinders were soaked only in buffer without any disinfectants.
167
Table 31 gives the determined ClO2 concentrations. Figure 51 shows the measured spectra in the
absence and presence of a penicylinder for a typical experiment.
Figure 51. Comparison of the ClO2 concentration change in the absence
(—) and in the presence (—) of a penicylinder. Initial concentrations:
[FAC] = 300 mg/L, [ClO2] = 300 mg/L, pH 7.0. See Table 31 for ClO2
concentrations at a given time.
Table 31. Comparison of the concentration change of
ClO2 in the absence and in the presence of a penicylinder.
Initial concentrations: [FAC] = 300 mg/L, [ClO2] =
300 mg/L, pH 7.0
[ClO2] mg/L
Time (minutes)
Penicylinder
absent
Penicylinder
present
30
62.8
78.8
120
0.5
5
From these figures in the presence of the penicylinders, the concentration of ClO2 is clearly
higher as compared with the absence of penicylinders. This means that the consumption of ClO2 is
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slower. Because of this, the concentration of the intermediates is possibly altered. However, based
only on the ClO2 concentration change it is hard to draw specific conclusions. Thus, determining the
concentration changes of the other species in the presence of the penicylinders is necessary. The
previously mentioned absorbance increase interferes with the spectrophotometric measurement of the
other species. Therefore, using iodometric titrations was necessary in further studies.
4.7.1. Results of iodometric measurements
The further kinetic studies were performed by following the changes in the concentrations of the
various species by iodometric titration. Iodometric measurements made it possible to determine the
concentration changes of all three species (FAC, ClO2, and chlorite ion) in the system. The effect of
the cylinders on the ClO2–FAC reaction was studied by using mixed disinfectant solutions. The initial
concentrations of the disinfectants were [FAC] = 300 mg/L, [ClO2] = 300 mg/L, and the pH was 7.5.
The following sections summarize the results of this study.
Figures 52 and 53 show the concentration change of FAC and ClO2. In these Figures, the
predicted values are based on the previously determined rate equation and rate constant. From these
Figures several conclusions can be drawn.
The concentration of ClO2 is significantly lower in these experiments than predicted by the rate
equation. The determined ClO2 concentrations are the lowest when no penicylinder was present in
the solution. The concentration of FAC is higher in the kinetic measurements than the predicted value
from the rate equation. In this case the highest FAC concentration is measured when no penicylinders
were present.
169
Figure 52. The concentration change of FAC in a FAC–ClO2
mixture. [FAC]0 = 300 mg/L, [ClO2]0 = 300 mg/L, pH = 7.0.
 predicted values, Ž in the absence of a penicylinder, • in
the presence of a penicylinder
Figure 53. The concentration change of ClO2 in a FAC–ClO2
mixture. [FAC]0 = 300 mg/L, [ClO2]0 = 300 mg/L, pH = 7.0.
 predicted values, Ž in the absence of a penicylinder, • in
the presence of a penicylinder
170
The difference in the measured and predicted ClO2 concentrations can be understood by
considering the large headspace over the solutions. Significant ClO2 loss is expected due to
evaporation from the solution. Thus the measured ClO2 concentrations are lower than predicted by
the rate equation. There is a difference between the measured ClO2 concentrations in the absence and
presence of penicylinders. This difference is similar to the measured concentration difference in the
spectrophotometric measurements. The measured ClO2 concentration is higher in the presence of the
cylinder than in the absence. This higher concentration is possibly due to a decrease in the evaporation
of ClO2 from the solution. The decrease in the evaporation can be the result of ClO2 being bound on
the cylinders, either in the pores of the porcelain cylinder or on the surface by forming complexes with
the proteins.
The concentration change of FAC can be interpreted based on the ClO2 concentration change.
It was found that the predicted ClO2 concentration is higher than the measured concentrations. The
lower than predicted ClO2 concentrations would result in a slower FAC–ClO2 reaction and in turn
less FAC is consumed. Therefore, the measured FAC concentrations are higher than the predicted
values. The difference in the FAC concentration in the presence and absence of the penicylinders can
be interpreted in a similar way. The measured ClO2 concentration is the lowest in the absence of any
penicylinders. The FAC–ClO2 is expected to be the slowest in this case and correspondingly the least
amount of FAC is consumed. Thus, the measured FAC concentration is the highest in the absence of
any penicylinders.
Further microbiological tests were conducted to prove the effect of ClO2 evaporation loss. In
these experiments the ClO2 loss was reduced by placing a glass wool plug above the solution to
minimize evaporation. The results were compared with control experiments which were conducted
171
under the same conditions, but no glass wool plug was used. The results showed that the number of
sterile penicylinders, when the ClO2 evaporation was minimized, was almost twice as high than in the
control experiments. During the AOAC testing phase, this means that by reducing the ClO2
evaporation loss it is possible to improve the “measured” the efficacy of the mixed disinfectant.
4.8. Conclusions
A new mixed disinfectant solution was developed which utilizes the reaction of FAC with ClO2.
The mixed disinfectant shows higher efficacy than the sum of the efficacy of the individual
disinfectants. This increased efficacy is due to the presence of reactive intermediates in the FAC–ClO2
mixture.
The microbiological experiments confirmed this improved efficacy. The mixed disinfectant
solution has a disinfection time of less than one hour. In comparison, the currently used FAC
disinfectant requires more than one day in order to achieve the same level of disinfection. In addition,
the dissolved chlorine concentration in the newly developed mixed disinfectant is only about one third
of the chlorine concentration currently used.
It has been shown that in the AOAC test procedure, the evaporation of ClO2 from the mixed
disinfectant solution can severely decrease the efficacy of the mixed disinfectant solution. A method
has been devised to minimize this evaporation and improve the efficacy of the mixed disinfectant.
The kinetic parameters of the FAC–ClO2 reaction in the disinfectant solution have been
determined. These parameters were used to interpret the results of the initial microbiological results.
The combination of the kinetic parameters and microbiological results was helpful in understanding
the deviations encountered during some of the microbiological measurements. By using the kinetic
172
and microbiological information the mixed disinfectant solution was improved with the use of a small
number of microbiological experiments. This resulted in significant savings in time and money.
Further work on the mixed disinfectant solution should include the development of an
electrochemical generator. This generator can make the mixed disinfectant solution a commercially
viable product.
173
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182
Appendix A
A.1 Program for collecting data from a Radiometer autotitrator
This program was written in Microsoft Quick BASIC. It is based on a simple terminal program,
which was provided as an example with the Quick BASIC package. The purpose of this program
is to receive the titration data from the Radiometer VIT 90 Videotitrator and save the data into a
text file. It is also possible to use the program to extract the titrant volums correponding to the
end points.
'Initialize the variables and the computer screen
DEFINT A-Z
DECLARE SUB Filter (InString$)
DIM lineIn$(500), num$(4)
num$(1) = “1.st“: num$(2) = “2.nd“: num$(3) = “3.rd“: num$(4) = “4.th“
Beginning:
COLOR 3, 1
CLS
Quit$ = CHR$(0) + CHR$(16)
'Ask for a file name to save the titration data. If no file name is entered
'the program is terminated. The default extension of the file is .tit. The
'program checks that the given file name is not longer than 8 characters.
inFile1:
INPUT "File name to save the raw data"; fileTitr$
IF fileTitr$ = "" THEN GOTO finish
IF LEN(fileTitr$) > 8 THEN
PRINT : PRINT "The file name cannot be longer than 8 characters."
PRINT "Try again! Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO inFile1
END IF
fileTitr$ = fileTitr$ + “.tit“
CLS
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'Set up prompt on bottom line of screen and turn cursor on. Printing on the
'screen between lines 1 and 23.
LOCATE 24, 1, 1
PRINT STRING$(80, "_");
LOCATE 25, 1
PRINT TAB(30); "Press ALT+q to quit";
VIEW PRINT 1 TO 23
'Open RS232 port
'9600 baud, no parity, 8-bit data, 1 stop bit, 256-byte input buffer
'Create the file to save the data.
OPEN "COM1:9600,N,8,1" FOR RANDOM AS #1 LEN = 256
OPEN fileTitr$ FOR APPEND AS #5
'This loop is to receive the data. The purpose of this loop is to "listen"
'at the open port and print the incoming data both on the screen and in the
'open file. It also checks if the exit key combination (ALT+q) is entered.
DO
KeyInput$ = INKEY
IF KeyInput$ = Quit$ THEN
EXIT DO
END IF
'Check the modem. If characters are waiting (EOF(1) is true), get them and
'print them to the screen and to the file.
IF NOT EOF(1) THEN
ModemInput$ = INPUT$(LOC(1), #1)
Filter ModemInput$
PRINT ModemInput$;
PRINT #5, ModemInput$;
END IF
LOOP
'If the exit key combination is entered, the port and the file are closed.
VIEW PRINT
CLOSE #1
CLOSE #5
CLS
'This part of the program asks the user if the inflection point data is to be
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'extracted from the titration file.
in1:
INPUT "Do you want to save the data points from the raw file"; dec$
IF dec$ = "n" OR dec$ = "N" THEN GOTO New
IF dec$ <> "y" AND dec$ <> "Y" THEN
PRINT : PRINT "Yes (y) or no (n)? Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO in1
END IF
'If the data is to be extracted, ask for the file name, comments, and the
'measured unit (mV or pH)
in2:
PRINT : INPUT "Do you want to add a comment"; dec$
IF dec$ = "n" OR dec$ = "N" THEN GOTO Processing
IF dec$ <> "y" AND dec$ <> "Y" THEN
PRINT : PRINT "Yes (y) or no(n)? Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO in2
END IF
PRINT : LINE INPUT "Comment: ", comment$
'This part extracts the data related to the end points and saves it into a
'separate file. The comment (if exists) is also added.
Processing:
counter% = 0
OPEN fileTitr$ FOR INPUT AS #1
DO UNTIL EOF(1)
LINE INPUT #1, lineIn$(counter%)
counter% = counter% + 1
LOOP
CLOSE #1
FOR i = 0 TO counter%
IF LEFT$(lineIn$(i), 2) = "#3" THEN
start% = i
ELSEIF LEFT$(lineIn$(i), 2) = "#4" THEN
final% = i - 1
END IF
NEXT i
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in3:
PRINT : INPUT "File name to save the titration data"; file$
IF file$ = "" THEN
PRINT : PRINT "File name? Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO in3
END IF
IF LEN(file$) > 8 OR file$ = "" THEN
PRINT : PRINT "The file name cannot be longer than 8 characters."
PRINT "Try again! Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO in3
END IF
file$ = file$ + ".dat"
in4:
PRINT "What is the measured unit?"
PRINT "a) pH"
PRINT "b) mV"
INPUT "Select the corresponding letter! ", unit$
IF unit$ <> "a" AND unit$ <> "A" AND unit$ <> "b" AND unit$ <> "B" THEN
PRINT : PRINT " What is the measured unit? Any key to continue."
WHILE INKEY$ = ""
WEND: CLS
GOTO in4
END IF
OPEN file$ FOR OUTPUT AS #1
PRINT #1, comment$: PRINT #1, : PRINT #1,
IF unit$ = "a" OR unit$ = "A" THEN
PRINT #1, "mL pH": unit$ = "pH"
ELSE PRINT #1, "mL mV": unit$ = "mL"
END IF
PRINT #1, MID$(lineIn$(start% + 1), 2, 11),
PRINT #1, MID$(lineIn$(start%), 4, 11)
FOR i = start% + 2 TO final% STEP 2
PRINT #1, MID$(lineIn$(i + 1), 2, 11),
PRINT #1, MID$(lineIn$(i), 2, 11)
NEXT i
PRINT #1, : PRINT #1,
PRINT #1, "Initial "; unit$; ": "
PRINT #1, MID$(lineIn$(final% + 1), 4, 11)
PRINT #1, : PRINT #1, "Inflection points "; unit$; " mL"
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FOR i = 1 TO 4
j = final% + (i - 1) * 3 + 3
PRINT #1, num$(i); " "; MID$(lineIn$(j), 2, 11); " "; MID$(lineIn$(j + 1),
2, 11)
NEXT i
CLOSE #1
New:
PRINT : INPUT "Do you want to run another titration (y/n)"; dec$
IF dec$ = "y" OR dec$ = "Y" THEN GOTO Beginning
finish:
END
'
'
'
'
'
========================= FILTER ==========================
Filters characters in an input string.
This subprogram is provided with the terminal program and
filters out line feed or backspace characters.
============================================================
SUB Filter (InString$) STATIC
' Look for backspace characters and recode them to
' CHR$(29) (the LEFT cursor key):
DO
BackSpace = INSTR(InString$, CHR$(8))
IF BackSpace THEN
MID$(InString$, BackSpace) = CHR$(29)
END IF
LOOP WHILE BackSpace
'Look for line-feed characters and remove any found:
DO
LineFeed = INSTR(InString$, CHR$(10))
IF LineFeed THEN
InString$ = LEFT$(InString$, LineFeed - 1) + MID$(InString$, LineFeed + 1)
END IF
LOOP WHILE LineFeed
END SUB
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A.2 Program for converting raw data from an Applied Photophysic SF to Excel format
This program was written to convert the raw data file from the SF into an Excel workbook. The
SF data file is a linear text file, which can not be directly used in Excel for data analysis.
'Initialize the variables used in the program.
Public Sub sfProcessing()
Dim fs, inFile(), numberOfFiles%, i%, currentFileNumber%, saveName$, timeRow%
Dim rwIndex%, lambdaRow%, dataRow%, topLeft, bottomRight, bias%, workFile
Dim rowNumber%, colNumber%, colIndex%, xCell, dataCollection
Dim workRange, nonEmpty%, col, tempBook
Const sigmaBlank = 0.005
'The program uses the built-in file search method of Visual Basic. It searches
'for files in the given directory based on the given file mask. The SF files
'have the ending of ,fff. The names of the found files are loaded into the
'the variable inFile, which is used later to open these files.
Set fs = Application.FileSearch
With fs
.LookIn = "c:\work\"
.Filename = "*,fff"
If .Execute > 0 Then
numberOfFiles = .FoundFiles.Count
ReDim inFile(numberOfFiles - 1)
For i = 0 To numberOfFiles - 1
inFile(i) = .FoundFiles(i + 1)
Next i
End If
End With
'The screen updating is turned off to increase the speed of the program.
Application.ScreenUpdating = False
'This loop opens the found SF files and converts them into Excel workbook.
currentFileNumber = -1
Do
currentFileNumber = currentFileNumber + 1
Workbooks.OpenText Filename:=inFile(currentFileNumber) _
, DataType:=xlDelimited, consecutivedelimiter:=True, Space:=True _
, fieldinfo:=Array(1, 1), trailingminusnumbers:=True
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saveName$ = ""
For i = 1 To Len(inFile(currentFileNumber))
If Mid$(inFile(currentFileNumber), i, 1) = "\" Then
saveName$ = ""
ElseIf Mid$(inFile(currentFileNumber), i, 1) = "," Then
Exit For
Else: saveName$ = saveName$ + Mid$(inFile(currentFileNumber), i, 1)
End If
Next i
'This part converts the SF file by moving the data and transposing it.
'The first part identifies the rows where the data blocks are located.
timeRow = 0: lambdaRow = 0: dataRow = 0
For rwIndex = 1 To 500
If Cells(rwIndex, 1) = "Times:" Then
timeRow = rwIndex
ElseIf Cells(rwIndex, 1) = "Lambda:" Then
lambdaRow = rwIndex
ElseIf Cells(rwIndex, 1) = "Data:" Then
dataRow = rwIndex
Exit For
End If
Next rwIndex
'Check if the file is valid
If timeRow = 0 Or lambdaRow = 0 Or dataRow = 0 Then GoTo skipFile:
Set topLeft = Cells(lambdaRow + 1, 1)
Set bottomRight = Cells(dataRow - 1, 1)
Range(topLeft, bottomRight).Copy
Cells(timeRow, 2).PasteSpecial Paste:=xlValues, Transpose:=True
Range(topLeft, bottomRight).Delete shift:=xlUp
Cells(lambdaRow + 1, 1).ClearContents
bias = 0
Do
bias = bias + 1
Cells(lambdaRow + 2, 1).Select
ActiveCell.CurrentRegion.Select
Selection.Copy
Cells(timeRow + bias, 2).PasteSpecial Paste:=xlValues, Transpose:=True
Cells(lambdaRow + 2, 1).Select
ActiveCell.CurrentRegion.Delete shift:=xlUp
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Cells(lambdaRow + 2, 1).Delete shift:=xlUp
Loop While Cells(lambdaRow + 2, 1) <> ""
Cells(lambdaRow, 1).ClearContents
For i = 1 To timeRow - 1
Rows(1).Delete
Next i
Cells(1, 1).Select
'This part saves the modified file as an Excel workbook in the same directory
'as the original SF file was found.
ActiveWorkbook.SaveAs "c:\work\" + saveName$ + ".xls", FileFormat:=xlNormal
ActiveWorkbook.Close
skipFile:
Loop Until currentFileNumber = numberOfFiles - 1
Application.ScreenUpdating = True
End Sub
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