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EFFECTS OF PHOSPHATE-BASED CORROSION
INHIBITORS ON DISINFECTANT STABILITY AND
HAA/NDMA FORMATION WHEN IN CONTACT WITH
COPPER, IRON AND LEAD
by
Hong Zhang
A thesis submitted in conformity with the requirements
for the degree of Doctor of Philosophy,
Graduate Department of Civil Engineering,
University of Toronto
© Copyright by Hong Zhang 2012
ii
EFFECTS OF PHOSPHATE-BASED CORROSION
INHIBITORS ON DISINFECTANT STABILITY AND
HAA/NDMA FORMATION WHEN IN CONTACT WITH
COPPER, IRON AND LEAD
Hong Zhang
Doctor of Philosophy, 2012
Graduate Department of Civil Engineering
University of Toronto
ABSTRACT
This research examined the impacts of water quality, phosphate-based corrosion
inhibitors and pipe wall exposure on free chlorine (HOCl)/chloramine (NH2Cl) degradation and
haloacetic acid (HAA)/N-nitrosodimethylamine (NDMA) formation in simulated distribution
system water mains and household plumbing at bench-scale and pilot scale.
In bench-scale bottle tests, the reactivity of fresh/pre-corroded pipe materials with
HOCl/NH2Cl in decreasing order was: ductile iron, copper, lead. The addition of phosphate-
based corrosion inhibitors generally increased HOCl/NH2Cl degradation for fresh iron coupons,
but decreased HOCl/NH2Cl decay only for fresh copper coupons. Generally, these corrosion
inhibitors did not impact HAA formation.
Copper corrosion products, including Cu(II), Cu2O, CuO and Cu2(OH)2CO3, catalyzed
HAA and NDMA formation. For HAAs, copper catalysis increased with increasing pH from 6.6
to 8.6 and/or increasing concentrations of these copper corrosion products. Interactions of copper
with natural organic matter (NOM), likely by complexation, and the subsequent increase in the
reactivity of NOM were proposed to be the primary reason for the increased HAA formation.
iii
NDMA formation increased with increasing Cu(II) concentrations, DMA concentrations,
alkalinity and hardness but was inhibited by the presence of NOM. The transformation of NH2Cl
to dichloramine (NHCl2) and complexation of copper with DMA were proposed to be involved
in elevating the formation of NDMA at pH 7.0.
Finally, in pilot-scale modified pipe loop tests, copper catalysis of NDMA formation was
confirmed, especially under laminar flow conditions, and iron was shown to possibly catalyze
NDMA formation under turbulent conditions. Orthophosphate increased the catalytic effects of
iron but decreased copper catalysis on NDMA formation by either modifying the properties of
the iron-associated suspended particles or reducing the dissolved metal concentrations.
Orthophosphate increased chloramine decay when in contact with iron, likely by promoting
nitrite formation, but orthophosphate decreased chloramine decay for copper and lead by
reducing the availability of metal corrosion products.
iv
ACKNOWLEDGMENTS
This research was supported financially by the Canadian Water Network, and the Natural
Sciences and Engineering Research Council of Canada.
I would like to thank my supervisor, Professor Susan Andrews, for her advice and
guidance on my research work. I would especially like to thank her for her patience and
encouragement throughout my studies, especially during my paper and thesis writing. I am
grateful for my committee members Professor Robert Andrews, Professor Brent Sleep, Professor
Ron Hofmann and Professor Michele Prevost for offering their suggestions. I also very much
appreciate Tim Walton from Municipal Region of Waterloo for providing me with many water
samples during my research study and generously assisting in the built-up of my pilot-scale pipe
loops. Help from Ian Douglas from City of Ottawa was also much appreciated. I would also
thank Srebri Petrov from the Department of Chemistry, University of Toronto and Peter
Broderson from the Department of Chemical Engineering, University of Toronto for performing
metal surface analyses and providing valuable suggestions for interpreting data. Special regards
are extended to Russell D’Souza in the Civil Engineering Environmental labs for his
professionalism and to Jennifer Lee and the student colleagues in Drinking Water Research
Group for their kind friendship.
Finally, I give special thanks to my husband (Jim) and my parents for their love and
encouragement.
v
TABLE OF CONTENTS
ABSTRACT .................................................................................................................................... ii
ACKNOWLEDGMENTS ............................................................................................................. iv
TABLE OF CONTENTS .................................................................................................................v
LIST OF ACRONYMS ...................................................................................................................x
LIST OF TABLES ....................................................................................................................... xiv
LIST OF FIGURES ..................................................................................................................... xvi
Introduction .................................................................................................................................1 1
1.1 Background ........................................................................................................................1
1.2 Objectives ...........................................................................................................................3
1.3 Thesis Organization ............................................................................................................5
1.4 References ..........................................................................................................................6
Literature Review ........................................................................................................................9 2
2.1 Disinfectants .......................................................................................................................9
Chemistry of Chlor(am)ination ................................................................................9 2.1.1
Disinfectant Residual Reactions in Distribution Systems .....................................11 2.1.2
Kinetic Models for Disinfectant Residual Decay in Distribution Systems............13 2.1.3
2.2 Disinfection Byproducts ...................................................................................................16
Haloacetic Acids ....................................................................................................16 2.2.1
N-nitrosodimethylamine ........................................................................................16 2.2.2
Disinfection By-Products in the Distribution System ............................................18 2.2.3
2.3 Corrosion ..........................................................................................................................20
Metal Dissolution and Precipitation in the Distribution System ...........................20 2.3.1
vi
Factors Affecting Metal Levels at the Tap ............................................................21 2.3.2
Corrosion Control ..................................................................................................23 2.3.3
2.4 Nitrification ......................................................................................................................24
2.5 Summary of Research Gaps .............................................................................................26
2.6 References ........................................................................................................................27
Free Chlorine Degradation and HAA Formation in Two Water Matrices in Contact with 3
Three Metal Materials ...............................................................................................................40
3.1 Introduction ......................................................................................................................42
3.2 Materials and Methods .....................................................................................................44
Reagents and Materials ..........................................................................................44 3.2.1
Experimental Procedures .......................................................................................45 3.2.2
3.3 Results and Discussion .....................................................................................................47
Free Chlorine Degradation .....................................................................................47 3.3.1
HAA Formation .....................................................................................................55 3.3.2
HAA Speciation .....................................................................................................62 3.3.3
3.4 Summary ..........................................................................................................................66
3.5 References ........................................................................................................................68
A Comparison of Iron, Copper and Lead Corrosion in Simulated Distribution Systems .........72 4
4.1 Introduction ......................................................................................................................74
4.2 Materials and Methods .....................................................................................................75
Reagents and Materials ..........................................................................................75 4.2.1
Experimental Procedures .......................................................................................77 4.2.2
Coupon Surface Analysis .......................................................................................78 4.2.3
4.3 Results and Discussion .....................................................................................................79
Metal Release and Phosphate-based Corrosion Inhibitors.....................................79 4.3.1
XPS Results for Corroded Coupons in Mannheim Water .....................................84 4.3.2
vii
Water Quality and Disinfectant Type ....................................................................89 4.3.3
4.4 Summary ..........................................................................................................................92
4.5 References ........................................................................................................................94
Catalytic Impacts of Copper Corrosion Products on Chlorine Decay and HAA Formation 5
in Simulated Distribution Systems ............................................................................................99
5.1 Introduction ....................................................................................................................101
5.2 Materials and Methods ...................................................................................................103
Reagents and Materials ........................................................................................103 5.2.1
Experimental Procedures .....................................................................................104 5.2.2
5.3 Results and Discussion ...................................................................................................105
Chlorine Decay ....................................................................................................106 5.3.1
HAA Formation and Speciation...........................................................................111 5.3.2
HAA Speciation ...................................................................................................115 5.3.3
5.4 Summary ........................................................................................................................117
5.5 References ......................................................................................................................118
Factors Affecting Copper Catalysis of NDMA Formation from DMA in Simulated 6
Premise Plumbing ...................................................................................................................122
6.1 Introduction ....................................................................................................................124
6.2 Materials and methods ....................................................................................................126
Chemicals and Materials ......................................................................................126 6.2.1
Experimental Procedures .....................................................................................127 6.2.2
6.3 Results and discussion ....................................................................................................128
Effect of Copper Concentrations .........................................................................128 6.3.1
Effects of DMA Concentration and Cu-NH2Cl Interactions ...............................129 6.3.2
Cu-DMA Complexation.......................................................................................131 6.3.3
Effect of Alkalinity ..............................................................................................132 6.3.4
Effect of NOM .....................................................................................................134 6.3.5
viii
Effect of Hardness................................................................................................135 6.3.6
Effect of pH and the Role of NHCl2 in NDMA Formation .................................136 6.3.7
6.4 Summary ........................................................................................................................138
6.5 References ......................................................................................................................139
Effects of Pipe Materials, Orthophosphate, and Flow Conditions on Chloramine Decay 7
and NDMA Formation in Modified Pipe Loops .....................................................................144
7.1 Introduction ....................................................................................................................146
7.2 Materials and Methods ...................................................................................................147
Reagents and Materials ........................................................................................147 7.2.1
Modified Pipe Loops............................................................................................148 7.2.2
Experimental Procedures .....................................................................................150 7.2.3
7.3 Results and Discussion ...................................................................................................151
Metal and Nitrogen Species Concentrations ........................................................151 7.3.1
Chloramine decay ................................................................................................154 7.3.2
NDMA formation.................................................................................................158 7.3.3
7.4 Summary ........................................................................................................................162
7.5 References ......................................................................................................................164
Conclusions .............................................................................................................................168 8
Practical Implications and Suggestions for Future Research ..................................................170 9
Appendices ..............................................................................................................................173 10
10.1 Authorisations to Include Copyright Material in Thesis ................................................173
10.2 QA/QC Protocols ...........................................................................................................177
GC/MS .................................................................................................................178 10.2.1
GC/ECD ...............................................................................................................180 10.2.2
Flame Atomic Absorption Spectrometry .............................................................185 10.2.3
Ion Chromatography ............................................................................................187 10.2.4
ix
10.3 Free Chlorine 24-hour Residuals through the Duration of Metal Coupon
Conditioning ...................................................................................................................190
10.4 HAA Speciation for Fresh Metal Coupons in Britannia Water ......................................191
10.5 XRD Analysis for Fresh Iron Coupons ..........................................................................192
10.6 Metal Release Kinetics and Results for Metal Surface Analysis ...................................193
10.7 NDMA Formation in the Presence of Orthophosphate in Modified Pipe Loops ...........197
10.8 Effects of Corrosion Inhibitors and the Extent of Metal Corrosion on
Monochloramine Degradation and NDMA Formation ..................................................198
Materials and Methods .........................................................................................199 10.8.1
Monochloramine Degradation .............................................................................202 10.8.2
NDMA Formation from DMA for Fresh Coupons ..............................................207 10.8.3
References ............................................................................................................212 10.8.4
10.9 Degradation Potential of Iron, Lead and Their Corrosion Products on HAA9 ...............216
Materials and Methods .........................................................................................217 10.9.1
HAA9 Degradation by Corroded Iron Coupons ...................................................218 10.9.2
HAA Degradation by Lead ..................................................................................223 10.9.3
References ............................................................................................................227 10.9.4
x
LIST OF ACRONYMS
AOB Ammonia-oxidizing bacteria
ACS Chemical purity designation as defined and published by the
American Chemical Society Committee on Analytical
Reagents
ANOVA Analysis of variance
APHA American Public Health Association
AWWA American Water Works Association
BCAA Bromochloroacetic acid
BDCAA Bromodichloracetic acid
BDOC Bioeliminable dissolved organic carbon
CDBAA Chlorodibromoacetic acid
CDPH California Department of Public Health
DBAA Dibromoacetic acid
DBP Disinfection byproducts
DCAA Dichloroacetic acid
D/DBPR Disinfectant/ Disinfection Byproduct Rule
DMA Dimethylamine
ECD Electron capture detector
FAAS Flame atomic absorption spectrometry
xi
GC/MS Gas chromatography/mass spectrometry
HAA Haloacetic acid
HOCl Hypochlorous acid
HPLC High-performance liquid chromatography
HRT Hydraulic retention time
H2SO4 Sulphuric acid
ID Inner diameter
LCR Lead and Copper Rule
LSD Least Significant Difference (Fisher’s)
MAC Maximum Acceptable Concentration
MCAA Monochloroacetic acid
MBAA Monobromoacetic acid
MCL Maximum Contaminant Level
MDL Method detection limit
MOE Ministry of the Environment (Ontario)
MS-FP Material-specific formation potential
MS-SDS Material-specific simulated distribution system
MWTP Mannheim Water Treatment Plant
NA Not available
NaOH Sodium hydroxide
xii
NCl3 Trichloramine or nitrogen trichloride
NDEA N-nitrosodiethylamine
NDMA N-nitrosodimethylamine
NDPA N-nitrosodi-n-propylamine
NDPhA N-nitrosodiphenylamine
NH2Cl Monochloramine
NHCl2 Dichloramine
NMOR N-nitrosomorpholine
NOB Nitrite-oxidizing bacteria
NOM Natural organic matter
NPIP N-nitrosopiperidine
NPYR N-nitrosopyrollidine
PACl Polyaluminum chloride
PVC Polyvinyl chloride
QA/QC Quality Assurance and Quality Control
SDS Simulated distribution system
SUVA UVA254/TOC
TBAA Tribromoacetic acid
TCAA Trichloroacetic acid
THM Trihalomethane
xiii
TOC Total organic carbon
UDMH Unsymmetrical dimethylhydrazine
USEPA United States Environmental Protection Agency
UV Ultraviolet
UVA254 Ultraviolet absorbance at a wavelength of 254nm
WEF Water Environment Foundation
XRD X-ray diffraction
XPS X-ray photoelectron spectroscopy
xiv
LIST OF TABLES
Table 2-1 Chlorine reactions in drinking water treatment ............................................................ 10
Table 2-2 Primary reactions of monochloramine ......................................................................... 11
Table 2-3 Free chlorine and chloramine model coefficients ........................................................ 15
Table 2-4 Relevant equilibrium reactions for copper in carbonated-bearing water ..................... 21
Table 3-1 Summary of water quality parameters for post-filtration water ................................... 45
Table 3-2 Summary of analytical methods ................................................................................... 46
Table 3-3 Chlorine decay constants for fresh metal coupons in the presence and absence of
corrosion inhibitors with Mannheim Water (n=4) ....................................................... 49
Table 3-4 Comparison of HOCl overall decay constants (h-1) for fresh coupons between two
water matrices (n=4) ..................................................................................................... 51
Table 3-5 Comparison of HOCl overall decay constants for corroded coupons between
Mannheim Water and Britannia Water (n=4) ............................................................... 54
Table 4-1 Summary of water quality parameters for two post-filtration water sources ............... 77
Table 4-2 Summary of analytical methods ................................................................................... 78
Table 5-1Water quality parameters for post-filtration water from MWTP ................................ 104
Table 7-1 Summary of water quality parameters for the influent of the pipe loops ................... 147
Table 7-2 Summary of design parameters for pipe loops ........................................................... 149
Table 7-3 Significance of the effects of flow conditions on chloramine decay determined by the
LSD test (95% confidence level) ................................................................................ 155
Table 7-4 Significance of the effects of orthophosphate on chloramine decay determined by the
LSD test (a confidence level of 95%) ........................................................................ 156
xv
Table 10-1 GC/MS method – Method detection limits for NDMA (n=8, 99% confidence level)
.................................................................................................................................... 178
Table 10-2 GC-ECD method – Method detection limits for HAA9 (n=8, 99% confidence level)
.................................................................................................................................... 181
Table 10-3 FAAS method – Method detection limits for iron, copper and lead ........................ 186
Table 10-4 IC method – Method detection limits for nitrite, nitrate, chloride, bromide, sulphate
and phosphate ............................................................................................................. 189
Table 10-5 Water quality parameters for post-filtration water ................................................... 200
Table 10-6 Summary of analytical methods ............................................................................... 201
Table 10-7 Comparison of NH2Cl decay constants for fresh coupons in two water matrices (n=4)
.................................................................................................................................... 202
Table 10-8 Comparison of NH2Cl overall decay constants for corroded coupons between
Mannheim Water and Britannia Water (n=4) ............................................................. 207
Table 10-9 Summary of the results for the single contrasts using the LSD test ......................... 209
xvi
LIST OF FIGURES
Figure 3-1 Free chlorine decay for fresh metal coupons with time in the presence and absence of
corrosion inhibitors in one set of experiments, n=2; (a) iron, (b) copper, (c) lead. ..... 48
Figure 3-2 Free chlorine overall decay constants for corroded coupons with Mannheim Water.
Initial free chlorine concentration 12.3 mg/L, error bars indicate standard deviation
(n=4) ............................................................................................................................. 52
Figure 3-3 Copper release kinetics for corroded coupons in Mannheim Water and Britannia
Water; error bars indicate the measured maximum and minimum values (n=2) ......... 55
Figure 3-4 Lead release kinetics for corroded coupons in Mannheim Water and Britannia Water;
error bars indicate the measured maximum and minimum values (n=2) ..................... 55
Figure 3-5 HAA6 formations with time in the presence and absence of corrosion inhibitors for
fresh metal coupons with Mannheim Water in one set of experiments, error bars
indicate the measured maximum and minimum values (n=2) ..................................... 56
Figure 3-6 HAA6 formations with time in the presence and absence of corrosion inhibitors for
fresh metal coupons with Britannia Water in one set of experiments, error bars indicate
the measured maximum and minimum values (n=2) ................................................... 59
Figure 3-7 HAA6 formation at 48 hours in the presence and absence of corrosion inhibitors for
corroded metal coupons with Mannheim Water and Britannia Water in one set of
experiments, error bars indicate the measured maximum and minimum values (n=2) 60
Figure 3-8 HAA formation and free chlorine demand for bulk water in the absence of corrosion
inhibitors ....................................................................................................................... 61
Figure 3-9 HAA formation and free chlorine demand for copper in the absence of corrosion
inhibitors ....................................................................................................................... 62
xvii
Figure 3-10 HAA speciation with time in the presence of 1 mg/L orthophosphate for fresh metal
coupons with Mannheim Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2) ......................................................... 63
Figure 3-11 HAA speciation with time in the presence of 1 mg/L orthophosphate for corroded
metal coupons with Mannheim Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2) ......................................................... 65
Figure 3-12 HAA speciation with time in the presence of 1 mg/L orthophosphate for corroded
metal coupons with Britannia Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2) ......................................................... 65
Figure 4-1 Kinetics of metal release from fresh metal coupons in the presence and absence of
corrosion inhibitors with HOCl in Mannheim Water; disinfectant concentrations, 12.3
mg/L; error bars indicate the measured maximum and minimum values (n=2) .......... 80
Figure 4-2 Comparison of XRD patterns for iron powders scratched from the surface of oxidized
iron coupons in the absence and presence of corrosion inhibitors ............................... 81
Figure 4-3 Comparison of XRD patterns for copper coupons in the absence and presence of
corrosion inhibitors. ...................................................................................................... 82
Figure 4-4 Comparison of XRD patterns for lead coupons in the absence and presence of 1 mg/L
orthophosphate ............................................................................................................. 82
Figure 4-5 Metal concentrations after 24 hours for corroded coupons in the absence/presence of
orthophosphate with HOCl and NH2Cl in Mannheim Water, initial disinfectant
concentrations 5.5 mg/L; error bars indicate the measured maximum and minimum
values (n=2) .................................................................................................................. 84
Figure 4-6 Comparison of elemental distribution for copper, lead and iron coupons in the
absence and presence of orthophosphate with NH2Cl ................................................. 86
Figure 4-7 Corrosion products of Fe, Cu and Pb as well as their relative distribution in the scales
...................................................................................................................................... 87
xviii
Figure 4-8 Comparison of iron concentrations at 24 hours in the presence of orthophosphate for
fresh and corroded coupons under free chlorine and chloramine, error bars indicate the
measured maximum and minimum values (n=2) ......................................................... 90
Figure 5-1 Chlorine degradation at different pH values in the absence and presence of 1 mg/L Cu
(II) for MWTP water; initial HOCl =10 mg/L; triplicate ........................................... 106
Figure 5-2 Chlorine decay for synthetic water in the presence and absence of 1 mg/L Cu(II) and
NOM; initial HOCl =4.2 mg/L; pH 8.3; duplicate ..................................................... 108
Figure 5-3 Chlorine degradation in the presence of different concentrations of Cu(II) for MWTP
water; initial HOCl =10 mg/L; triplicate; pH 8.3 ....................................................... 109
Figure 5-4 Pseudo-first-order decay rates of free chlorine for MWTP water containing dissolved
Cu(II) and solid copper corrosion products ................................................................ 110
Figure 5-5 HAA9 at 100 hours at different pH values in the absence and presence of 1 mg/L
Cu(II) for MWTP water, initial HOCl =10 mg/L (error bars represent standard
deviation of triplicate tests) ........................................................................................ 112
Figure 5-6 HAA9 formation in the presence of Cu(II) with varying concentrations at pH 8.3 for
MWTP water (error bars represent standard deviation of triplicate tests) ................. 113
Figure 5-7 HAA9 at a reaction time of 32 hours in the presence of copper corrosion solids at pH
8.3 for MWTP water (error bars represent standard deviation of triplicate tests) ...... 114
Figure 5-8 HAA9 formation rates for MWTP water containing dissolved Cu(II) without solids
and solid copper corrosion products ........................................................................... 115
Figure 5-9 Effect of pH on HAA speciation in the absence and presence of 1 mg/L Cu(II) at
reaction time of 100 hours for MWTP water (error bars represent standard deviation of
triplicate tests) ............................................................................................................ 116
Figure 6-1 NDMA formation with increasing copper concentrations from added CuSO4; pH 7.0,
initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA spiked, 24 hours, error bars indicate the
measured maximum and minimum values (n=2) ....................................................... 129
xix
Figure 6-2 NDMA formation kinetics in the absence and presence of copper; pH 7, Cu(II) 1
mg/L spiked, initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA, error bars indicate the
measured maximum and minimum values (n=2) ....................................................... 129
Figure 6-3 NDMA formation with DMA concentrations; pH 7.0, initial NH2Cl 2.3±0.1 mg/L,
Cu(II) 1 mg/L spiked, 24 hours, d error bars indicate the measured maximum and
minimum values (n=2) ............................................................................................... 130
Figure 6-4 Chloramine speciation in the absence and presence of copper at pH 6.7 and 7, Milli-Q
water ........................................................................................................................... 131
Figure 6-5 Dissolved copper concentrations with varying DMA; filtered by 0.2 µm Nylaflo®
Nylon membrane filter paper, pH 7, Cu(II) 1 mg/L spiked, 6 hours, error bars indicate
the measured maximum and minimum values (n=2) ................................................. 132
Figure 6-6 NDMA formation with increasing alkalinity and copper speciation as a function of
alkalinity (determined by MINEQL+ version 4.5); pH 7.0, initial NH2Cl 2.3±0.1
mg/L, Cu(II) 1 mg/L spiked, 11.2 µg/L DMA, 24 hours, duplicate ......................... 133
Figure 6-7 NDMA, NH2Cl residual and dissolved copper with increasing SR-NOM
concentrations; pH 7.0, initial NH2Cl 2.3±0.1 mg/L, Cu(II) 1 mg/L spiked, 11.2 µg/L
DMA, 24 hours, error bars indicate the measured maximum and minimum values
(n=2) ........................................................................................................................... 134
Figure 6-8 Variation of NDMA formation and dissolved copper concentrations with increasing
hardness; pH 7.0, Cu(II) 1 mg/L spiked, initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA
spiked, SR-NOM 4.1 mg/L as TOC, 24 hours, error bars indicate the measured
maximum and minimum values (n=2) ....................................................................... 136
Figure 6-9 NDMA formation at four pH levels in the absence and presence of 1 mg/L Cu(II) in
Milli-Q water; 24 hour, NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA, error bars indicate the
measured maximum and minimum values (n=2) ....................................................... 137
Figure 7-1 Module configuration (adapted from Cantor, 2009) ................................................. 148
Figure 7-2 Schematic of the modified pipe loop (not to scale)................................................... 149
xx
Figure 7-3 Metal concentrations in the pipes of iron, copper and lead in the absence and presence
of 1 mg/L orthophosphate under two flow conditions ............................................... 152
Figure 7-4 Variations of nitrate, nitrite and ammonia in the lead loop....................................... 153
Figure 7-5 Pseudo-first-order chloramine decay constants for the four pipe loops under different
flow conditions, initial chloramine 1.3~1.6 mg/L as Cl2, error bars indicate the
measured maximum and minimum values (n=2) ....................................................... 154
Figure 7-6 Box and Whisker plots for nitrite concentrations before and after the addition of
orthophosphate overlaid with pseudo-first-order chloramine decay rate constants,
laminar flow (n=5), turbulent flow (n=6). The top and bottom of the box represent the
75th and 25th percentile, respectively, while the whiskers represent the maximum and
minimum values ......................................................................................................... 157
Figure 7-7 NDMA formation in four pipe loops in the absence of orthophosphate under two flow
conditions ................................................................................................................... 159
Figure 10-1 Example of GC/MS calibration curve for NDMA (NDMA 0~200 ng/L, d6-NDMA
50 ng/L) ...................................................................................................................... 178
Figure 10-2 GC/MS method for NDMA- spike recovery chart .................................................. 179
Figure 10-3 Example of GC-ECD calibration curves for HAA9 (HAA9 0~100 µg/L, 2,3,5,6-
TFBA 100 µg/L) ......................................................................................................... 180
Figure 10-4 Example of GC-ECD calibration curves for HAA9 (HAA9 0~100 µg/L, 2,3,5,6-
TFBA 100 µg/L), continued ....................................................................................... 181
Figure 10-5 GC-ECD method for MCAA- spike recovery chart ............................................... 182
Figure 10-6 GC-ECD method for MBAA- spike recovery chart ............................................... 182
Figure 10-7 GC-ECD method for DCAA- spike recovery chart ................................................ 183
Figure 10-8 GC-ECD method for TCAA- spike recovery chart................................................. 183
xxi
Figure 10-9 GC-ECD method for DBAA- spike recovery chart ................................................ 184
Figure 10-10 GC-ECD method for BCAA- spike recovery chart .............................................. 184
Figure 10-11 Example of FAAS calibration curves for iron, copper and lead ........................... 185
Figure 10-12 Example of IC calibration curves for chloride, sulphate and bromide ................. 187
Figure 10-13 Example of IC calibration curves for nitrite, nitrate and phosphate ..................... 188
Figure 10-14 HAA speciation with time in the presence 1 mg/L orthophosphate for fresh metal
coupons with Britannia Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2) ....................................................... 191
Figure 10-15 Comparison of XRD patterns for oxidized iron coupon in the presence of
polyphosphate-orthophosphate blends and polished iron coupon as control. ............ 192
Figure 10-16 Kinetics of metal release from fresh metal coupons in the presence and absence of
corrosion inhibitors with NH2Cl in Mannheim Water; disinfectant concentrations, 12.3
mg/L; error bars indicate the measured maximum and minimum values (n=2) ........ 193
Figure 10-17 Kinetics of metal release from fresh metal coupons with time in the presence and
absence of corrosion inhibitors in Britannia Water; disinfectant concentrations, 12.3
mg/L; error bars indicate the measured maximum and minimum values (n=2) ........ 194
Figure 10-18 Kinetics of metal release from corroded coupons in the absence/presence of
orthophosphate in Britannia Water, initial disinfectant concentrations 5.5 mg/L; error
bars indicate the measured maximum and minimum values (n=2) ............................ 195
Figure 10-19 Comparison of elemental distribution for iron, copper and lead coupons in the
absence and presence of orthophosphate with HOCl ................................................. 196
Figure 10-20 NDMA formation in four pipe loops in the presence of orthophosphate under two
flow conditions ........................................................................................................... 197
Figure 10-21 NH2Cl overall decay constants for fresh coupons with Mannheim Water; NH2Cl
12.3 mg/L, error bars indicate the measured maximum and minimum values (n=2) 203
xxii
Figure 10-22 Released metal concentrations at 24 hours for fresh metal coupons in Mannheim
Water; NH2Cl 12.3 mg/L; error bars indicate the measured maximum and minimum
values (n=2) ................................................................................................................ 204
Figure 10-23 Released iron and copper concentrations at 24 hours in the absence and presence of
corrosion inhibitors for fresh metal coupons in Britannia Water; NH2Cl 12.3 mg/L;
error bars indicate the measured maximum and minimum values (n=2) ................... 206
Figure 10-24 NDMA formation from DMA as a result interactions of metal coupons, corrosion
inhibitors and NH2Cl. Initial NH2Cl 14.5±0.5 mg/L, DMA 1 µM/L, pH 8.3, in Milli-
Q, error bars indicate standard deviation (n=3) .......................................................... 208
Figure 10-25 Total iron concentrations and NDMA formation for iron coupons in the absence
and presence of corrosion inhibitors, error bars indicate standard deviation (n=3) ... 210
Figure 10-26 Cu(II) concentrations and NDMA formation for copper coupons in the absence and
presence of corrosion inhibitors, error bars indicate standard deviation (n=3) .......... 211
Figure 10-27 Reduction of HAA9 by corroded iron coupons in the absence and presence of
different corrosion inhibitors. Milli-Q water, error bars indicate the measured
maximum and minimum values (n=2) ....................................................................... 219
Figure 10-28 HAA degradation by corroded iron coupons treated with poly/ortho blends, Milli-Q
water, error bars indicate the measured maximum and minimum values (n=2) ........ 219
Figure 10-29 Reduction of MCAA by corroded iron coupons in the absence and presence of
different corrosion inhibitors. Milli-Q water, error bars indicate the measured
maximum and minimum values (n=2) ....................................................................... 220
Figure 10-30 Comparison of XRD patterns for oxidized iron coupon in the presence of
poly/ortho-phosphate blends and polished iron coupon as control. Red: iron coupon as
control; Black: non-scratched iron coupon with poly/ortho-phosphate blends .......... 220
Figure 10-31 HAA speciation degradation by corroded iron coupons treated with poly/ortho
blends, error bars indicate the measured maximum and minimum values (n=2)....... 221
xxiii
Figure 10-32 Reduction of HAA9 by fresh lead coupons, Milli-Q water, pH 8.3, error bars
indicate the measured maximum and minimum values (n=2) ................................... 224
Figure 10-33 Degradation of MCAA by fresh lead coupons, Milli-Q water, pH 8.3, error bars
indicate the measured maximum and minimum values (n=2) ................................... 224
Figure 10-34 Degradation of single HAA species by fresh lead coupons, Milli-Q water, pH 8.3,
error bars indicate the measured maximum and minimum values (n=2) ................... 225
Figure 10-35 Degradation of HAA species in the presence of 1g/L lead corrosion products. Milli-
Q water, pH 8.3, error bars indicate the measured maximum and minimum values
(n=2) ........................................................................................................................... 227
1
Introduction 1
1.1 Background
Free or combined chlorine is commonly applied as a disinfectant residual to prevent
water quality deterioration in distribution systems beyond the treatment plant and to ensure the
delivery of safe and clean water to consumers. However, there are drawbacks to chemical
disinfection. It increases the risks of disinfection by-product (DBP) formation when these
secondary disinfectants react with natural organic matter present in the water. During
chlorination, trihalomethanes (THMs) and haloacetic acids (HAA9) are the dominant classes of
chlorinated DBPs, whereas combined chlorine potentially increases the formation of
nitrosamines (Choi et al., 2002; Najm and Trussell, 2001). This thesis focuses on HAAs and
nitrosamines because THM regulations are already well-established, while HAA and nitrosamine
regulations have been under recent development. In addition, switching from free chlorine to
combined chlorine to decrease HAAs (and THMs) has been observed to increase nitrosamines.
Therefore, nitrosamine formation might be an unintended consequence of meeting new HAA
regulations.
These DBPs have potential health risks. For example, the toxicological properties of
HAAs vary with the specific compound. Dichloroacetic acid (DCAA) is considered to be a
probable carcinogen, and trichloroacetic acid (TCAA) is regarded as a possible carcinogen to
humans (Health Canada, 2008). The Stage 2 Disinfectant/ Disinfection Byproduct Rule
(D/DBPR) regulates the Maximum Contaminant Level (MCL) of HAA5 (monochloroacetic acid,
DCAA, TCAA, monobromoacetic acid, and dibromoacetic acid) at 60 µg/L (USEPA, 2006). The
Guidelines for Canadian Drinking Water Quality established the Maximum Acceptable
Concentration (MAC) for HAA5 in drinking water at 80 µg/L based on a locational running
annual average of a minimum of quarterly samples taken from the distribution system (Health
Canada, 2008). N-nitrosodimethylamine (NDMA) is classified as a probable human carcinogen,
and guidelines to limit its concentration in drinking water vary. Health Canada has established
the guideline for NDMA in drinking water at a MAC of 40 ng/L (Health Canada, 2011). The
Ontario Ministry of the Environment, Canada, has set a MAC for NDMA of 9 ng/L (MOE,
2
2002). The California Department of Public Health has set a notification level of 10 ng/L for
NDMA (CDPH, 2009).
As potential oxidants, secondary disinfectants may also increase metal corrosion rates for
distribution pipes and household plumbing materials. Although confounded by other factors such
as pH, alkalinity and pipe age, free chlorine is generally more aggressive than combined chlorine
in metal corrosion due to its higher oxidation potential. Increased iron and copper corrosion rates
have been observed during chlorination (Boulay and Edwards, 2001; Cantor et al., 2003;
MacQuarrie et al., 1997; Rahman et al., 2007). However, increased lead leaching has been linked
to the switch of the disinfectant residual from free chlorine to monochloramine (Renner, 2004).
Despite some disadvantages associated with combined chlorine, the application of combined
chlorine will receive increasing attention due to the implementation of Stage 2
Disinfectant/Disinfection Byproduct Rule (D/DBPR) (USEPA, 2006) and the need to decrease
chlorinated DBPs.
To date, extensive studies have investigated disinfectant decay and corrosion processes
for lead, copper and iron individually in simulated and/or full distribution systems where free
chlorine and chloramine are used (including but not limited to Rossman et al., 2001; Haas et al.,
2002; Hallam et al., 2002; Al-Jasser, 2007; Lytle and Schock, 2005; McNeill and Edwards,
2001; Sarin et al., 2003; Xiao et al., 2007). These studies have examined the roles of physical
and chemical compositions of water, biofilms, and hydrodynamics on disinfectant residual
stability and corrosion. The addition of phosphate-based corrosion inhibitors as a corrosion
control strategy has been widely employed by many water systems to achieve compliance with
the USEPA’s Lead and Copper Rule (LCR) (USEPA, 2008). These corrosion inhibitors may
interfere with metal solubility and speciation, thereby making interactions between corrosion
products, disinfectant residual and DBP formation more complicated. To the author’s
knowledge, to date, no research has been conducted to investigate the interactive impacts of
disinfectant residual, metal corrosion and corrosion inhibitors on DBP formation (especially for
HAAs or nitrosamines such as NDMA). Furthermore, there have been no reports that compare
the behavior of disinfectant degradation and DBP formation for all three metal materials in the
same water matrix to develop a more comprehensive understanding about the impacts of metal
materials and their interactions with corrosion inhibitors on disinfectant decay and DBP
formation.
3
Therefore, several questions that could be asked include:
Are there differences in disinfectant degradation and HAA/NDMA formation for
different pipe materials treated with different phosphate-based corrosion
inhibitors?
Do metal age and water quality affect the efficacy of corrosion inhibitors and
disinfectant decay as well as HAA/NDMA formation?
How do flow conditions potentially affect secondary disinfectant stability and
NDMA formation?
What are the impacts of metal corrosion products on the fate of HAA/NDMA in
distribution systems?
An improved understanding of these interactions is needed as a basis for the development of
effective mitigation strategies.
1.2 Objectives
The overall objective of this research was to evaluate the impacts that water quality
characteristics (e.g. pH and alkalinity), phosphate-based corrosion inhibitors, pipe wall exposure
and flow conditions may have on disinfectant residual stability and the formation of disinfection
byproducts (e.g., haloacetic acids and nitrosamines) in distribution systems. Regulatory
guidelines for HAA5 and N-nitrosodimethylamine (NDMA) have recently been added to the
“Guidelines for Canadian Drinking Water Quality” (Health Canada, 2008 and 2011). This
information, therefore, has increased the awareness of water utilities and households to minimize
HAAs and NDMA in their water systems. Given that the regulations for THMs have been
already well-established in Canada, their formation in distribution systems is, therefore, beyond
the scope of this research. Results of this research will advance the understanding of the complex
physical, chemical and, to some extent, biological reactions which potentially cause disinfectant
degradation and HAA/NDMA formation in distribution system water mains and household
plumbing. The specific objectives of this thesis were to:
1) Investigate disinfectant degradation and HAA/NDMA formation in different water matrices
under the influence of metal corrosion and corrosion inhibitors. No relevant information has
been found in the literature regarding the impacts of corrosion inhibitors and/or their
4
interactions with metal surface on disinfectant degradation and HAA/NDMA formation in
newly installed or the aged pipe systems. Therefore, there is a need to investigate the
impacts of metal corrosion and corrosion preventive strategies on: a) disinfectant residual
stability and b) the fate of HAA/NDMA in distribution mains and household plumbing.
2) Examine the impact of disinfectant types, corrosion inhibitors and water quality on metal
release kinetics using material-specific simulated distribution system tests. Metal corrosion
is a primary reason for disinfectant residual degradation in distribution systems. To date,
most studies investigating metal corrosion processes have been conducted by performing
either pilot-scale pipe rig (loop) tests or full-scale distribution system tests. Impacts of
corrosion inhibitors, disinfectant types and water quality on metal corrosion may have been
confounded by hydrodynamics and microorganisms, and thus cannot be fully elucidated.
Therefore, the investigation of factors influencing metal corrosion under controlled
experimental conditions is important for evaluating the impacts of metal types, corrosion
inhibitors and water quality on disinfectant degradation and HAA/NDMA formation.
3) Evaluate the catalytic potential of copper and its corrosion products during HAA9 and
NDMA formation. The investigation of copper catalysis during HAA9 and NDMA formation
is essential as it may affect the fate of HAA9 and NDMA in household plumbing where
copper pipes and alloys are widely used. Limited research to date has been performed to
look into copper catalytic potential, the factors affecting copper catalysis (including pH, the
presence of organic matter, and copper concentrations), and the likely mechanisms of copper
catalysis during HAA9 and NDMA formation.
4) Examine the effects of pipe materials and orthophosphate on chloramine decay and NDMA
formation under different flow conditions using modified pipe loops. Information about
chloramine degradation kinetics and NDMA formation under the influence of pipe materials,
corrosion inhibitors, nitrite formation, and hydrodynamic conditions is limited and
inconsistent. This demonstrates a need to identify the factors which may play a key role in
maintaining chloramine residual and the formation of NDMA in pipes of different metallic
materials.
5
1.3 Thesis Organization
This thesis is composed of a general introduction (Chapter 1), a brief literature review
(Chapter 2), and five results chapters. Early experiments to survey the possible impacts of free
chlorine on HAA formation and corrosion are summarized and separated into Chapters 3 and 4.
Chapter 3 provides detailed discussion about the effects of corrosion inhibitors and their
interactions with metal surfaces on free chlorine degradation and HAA formation. Chapter 4
examines the metal release kinetics of iron, copper and lead through the use of these bench-scale
material-specific simulated distribution system tests and material-specific formation potential
tests. Copper catalysis of HAA and NDMA formation, under controlled bench-scale
experimental conditions, are reported in Chapters 5 and 6, respectively. Chapter 7 looks further
into some NDMA formation issues and discusses the results of pilot-scale tests regarding the
impacts of pipe materials, orthophosphate and also flow conditions on chloramine decay and
NDMA formation. Additional tests to investigate monochloramine degradation and NDMA
formation as well as the potential of iron, lead and their corrosion products to degrade HAA
results were initiated but not completed to the same extent as the other chapters, and thus the
results are presented in the Appendices 10.8 and 10.9, respectively.
In general, each results chapter is a separate manuscript that has been submitted for peer-
reviewed journal publication. At the time of writing, some contents from Chapters 3, 4 and 5
have been published as:
Zhang, H., Andrews, S. A. (2012) Effects of Phosphate-based Corrosion
Inhibitors on the Kinetics of Chlorine Degradation and HAA Formation in
Contact with Three Metal Materials. Canadian Journal of Civil Engineering, 39,
44-54.
Zhang, H., Andrews, S.A. (2012) Catalysis of Copper Corrosion Products on
Chlorine Decay and HAA Formation in Simulated Distribution Systems. Water
Research, 46 (8), 2665-2673.
The contents from Chapters 6 and 7 have been submitted as:
Zhang, H., Andrews, S.A. Factors Affecting Copper Catalysis during NDMA
Formation from DMA in Simulated Premise Plumbing. Water Research.
6
Zhang, H., Andrews, S.A. Effects of Pipe Materials, Orthophosphate, and Flow
Conditions on Chloramine Decay and NDMA Formation in Modified Pipe Loops.
Journal of Water Supply: Research and Technology – Aqua.
In addition, a separate Methods and Materials chapter is not included in this thesis
because details of the methods and materials employed in each experiment have been explained
in the relevant results chapters/papers. Similarly, while a brief literature review is included as
Chapter 2 to provide an overall context for the thesis, the papers in the results chapters contain
additional references as appropriate. QA/QC protocols followed in all experiments are
summarized in Chapter 10 (Appendices 10.2).
All the work and writing presented in this thesis were done by the author, with review
and editing by Professor Susan Andrews. Metal coupon surface analyses (Chapters 4 and 7) were
performed by Srebri Petrov from the Department of Chemistry, University of Toronto and Peter
Broderson from the Department of Chemical Engineering, University of Toronto.
1.4 References
Al-Jasser, A.O. (2007) Chlorine decay in drinking-water transmission and distribution systems:
Pipe service age effect. Water Research, 41(2), 387-396.
Boulay, N., and Edwards, M. (2001) Role of temperature, chlorine, and organic matter in copper
corrosion by-product release in soft water. Water Research, 35(3), 683-690.
California Department of Public Health (CDPH). California Drinking Water: NDMA and Other
Nitrosamines - Drinking Water Issues. http://www.cdph.ca.gov/certlic/drinkingwater/
Pages/NDMA.aspx (accessed September 15, 2011).
Cantor, A. F., Park, J. K., and Vaiyavatjamai, P. (2003) Effect of chlorine on corrosion in
drinking water systems. Journal American Water Works Association, 95(5), 112-123.
Choi, J., Duirk, S. E., and Valentine, R. L. (2002) Mechanistic studies of N-
nitrosodimethylamine (NDMA) formation in chlorinated drinking water. Journal of
Environmental Monitoring, 4(2), 249-252.
7
Haas, C.N., Gupta, M., Chitluru, R. and Burlingame, G. (2002) Chlorine demand in disinfecting
water mains. Journal of American Water Works Association, 94(1), 97.
Hallam, N.B., West, J.R., Forster, C.F., Powell, J.C. and Spencer, I. (2002) The decay of chlorine
associated with the pipe wall in water distribution systems. Water Research, 36(14),
3479-3488.
Health Canada (2008) Guidelines for Canadian Drinking Water Quality: Guideline Technical
Document — Haloacetic Acids., Water, Air and Climate Change Bureau, Healthy
Environments and Consumer Safety Branch, Health Canada, Ottawa, Ontario.
Health Canada (2011). Guidelines for Canadian Drinking Water Quality: Guideline Technical
Document- N-Nitrosodimethylamine (NDMA). Water, Air and Climate Change Bureau,
Healthy Environments and Consumer Safety Branch, Health Canada, Ottawa, Ontario.
(Catalogue No H128-1/11-662E).
Lytle, D. A., Sarin, P., and Snoeyink, V. L. (2005) The effect of chloride and orthophosphate on
the release of iron from a cast iron pipe section. Journal of Water Supply Research and
Technology-Aqua, 54(5), 267-281.
McNeill, L.S. and Edwards, M. (2001) Iron pipe corrosion in distribution systems. Journal of
American Water Works Association, 93(7), 88-100.
MacQuarrie, D. M., Mavinic, D. S., and Neden, D. G. (1997) Greater Vancouver Water District
drinking water corrosion inhibitor testing. Canadian Journal of Civil Engineering, 24(1),
34-52.
Ministry of the Environment (MOE). Safe Drinking Water Act 2002. Ontario Regulation 169/03,
Schedule 2. http://www.e-laws.gov.on.ca/html/regs/english/elaws_regs_030169_e.htm
(accessed September 15, 2011)
Najm, I., and Trussell, R. R. (2001) NDMA formation in water and wastewater. Journal
American Water Works Association, 93(2), 92-99.
8
Rahman, S., McDonald, B. C., and Gagnon, G. A. (2007) Impact of secondary disinfectants on
copper corrosion under stagnation conditions. Journal of Environmental Engineering-
Asce, 133(2), 180-185.
Renner, R. (2004) Plumbing the depths of DC's drinking water crisis. Environmental Science &
Technology, 38(12), 224A-227A.
Rossman, L. A., Brown, R. A., Singer, P. C., and Nuckols, J. R. (2001) DBP formation kinetics
in a simulated distribution system. Water Research, 35(14), 3483-3489.
Sarin, P., Clement, J.A., Snoeyink, V.L. and Kriven, W.W. (2003) Iron release from corroded
unlined cast-iron pipe. Journal American Water Works Association, 95(11), 85.
USEPA (2006) 40 CFR Parts 9, 141, and 142 National Primary Drinking Water Regulations:
Stage 2 Disinfectants and Disinfection Byproducts Rule. Federal Register 71(2), 387-493.
USEPA (2008) Lead and Copper Rule: A Revised Quick Reference Guide. http://water.epa.gov/
lawsregs/rulesregs/sdwa/lcr/upload/LeadandCopperQuickReferenceGuide_2008.pdf
(accessed October, 2011).
Xiao, W.Z., Hong, S.K., Tang, Z.J. and Taylor, J.S. (2007) Effects of blending on total copper
release in distribution systems. Journal American Water Works Association, 99(1), 78-88
9
Literature Review 2
To date, considerable efforts have been made to address the protection of water quality in
distribution systems, mainly by optimizing in-plant treatment processes. Strategies to stabilize
disinfectant residual while simultaneously controlling DBP formation include water quality
adjustment and DBP precursor removal. Fewer studies have been conducted on the fate of DBPs,
especially haloacetic acids (HAA9) and nitrosamines, in distribution systems, under the
interactive impacts of disinfectant residual and the internal pipe environment. This literature
review begins with a brief review of disinfectant chemistry and its application in distribution
systems. The mechanisms responsible for HAA9 and nitrosamine (especially N-
nitrosodimethylamine, NDMA) formation and influential factors associated with water quality
and distribution system characteristics are summarized. Finally, other aspects of water quality
deterioration in distribution systems resulting from disinfectant application are reviewed
(corrosion and nitrification).
2.1 Disinfectants
Chemistry of Chlor(am)ination 2.1.1
Free and combined chlorine are commonly used as secondary disinfectants to reduce the
possible occurrence of biological regrowth. The dominance of free chlorine species (HOCl, OCl-
and Cl2) is dependent on pH. Under the typical pH range in distribution systems, hypochlorous
acid (HOCl) and hypochlorite ion (OCl-) are the main chlorine species. Normally, HOCl is
predominant at pH lower than 7.5 (the pKa of HOCl), whereas OCl- is the main species at higher
pH (> 7.5). Reactions that free chlorine can undergo in drinking water treatment generally
involve oxidation and substitution. The significant reactions of chlorine are summarized in Table
2-1.
10
Table 2-1 Chlorine reactions in drinking water treatment
Reaction type Examples
Ammonia substitution NH3+HOClNH2Cl+H2O
NH2Cl+HOClNHCl2+H2O
Inorganic oxidation
Mn2+
+HOCl+2H2O→MnO(OH)2+3H++Cl
-
2Fe2+
+HOCl+5H2O→2Fe(OH)3+5H++Cl
-
NO2-+HOCl+2H2O→NO3
-+H
++Cl
-
HSO3-+HOCl→SO4
2-+2H
++Cl
-
HS-+HOCl→S+2H
++H2O +Cl
-
Br-+HOCl→HOBr +Cl
-
Organic reaction oxidation RCHO+HOCl→RCOOH +H
++Cl
-
RCOOH+HOCl→R+CO2+ H++Cl
-
Decomposition 3HOCl→3H
++2Cl
-+ClO3
-
2HOCl→2H++2Cl
-+O2
(Source: Dickenson, 2005)
In the presence of ammonia, free chlorine reacts rapidly and in a stepwise manner to form
chloramine compounds, including monochloramine (NH2Cl), dichloramine (NHCl2), and
trichloramine (NCl3). The application of NH2Cl has received increasing attention due to
requirements in the U.S. to comply with the Stage 2 Disinfectant/Disinfection Byproducts Rule
(D/DBPR) and the need to decrease chlorinated DBPs. NHCl2 and NCl3 formation in drinking
water are avoided due to odor and taste problems. The major reactions with regard to the
complex chemistry of chloramines have been summarized in Table 2-2.
Usually, NH2Cl is the dominant species of chloramines. Its rapid formation is observed at
a pH value of approximately pH 8.0 and a Cl2/N mass ratio of less than 5:1, which represents
common conditions for drinking water chloramination. A Cl2/N mass ratio from 3:1 to 5:1,
typically 4:1, is currently used to form NH2Cl to minimize the risks of NHCl2 and DBP
formation while reducing nitrification and biofilm growth associated with the excess of NH3
(USEPA, 1999). NHCl2 is formed at higher Cl2/N ratios (mass ratio >5:1) or at low pH values
(pH <5). Its formation can also be catalyzed by the presence of acetic acid, phosphate, carbonate,
and silicate species (Schreiber and Mitch, 2007; Trofe et al., 1980; Valentine et al., 1988;
Vikesland et al., 2001).
11
Table 2-2 Primary reactions of monochloramine
Reaction Stoichiometry Rate/Equilibrium
Constant (25 ºC) Reaction Stoichiometry
Rate/Equilibrium
Constant (25 ºC)
HOCl+NH3→NH2Cl+H2O k1=1.5X1010
M-1
h-1 b
I+ NHCl2→HOCl+ N2+ HCl k9=1.0 X 108M
-1h
-1
NH2Cl+H2O→HOCl+NH3 k2=7.6 X 10-2
h-1 b
I+ NH2Cl→N2+ H2O +HCl k10=3.0 X 107M
-1h
-1
HOCl+ NH2Cl→NHCl2+ H2O k3=1.0 X 106M
-1h
-1 NH2Cl+H+NH3Cl
+ k11=28M-1
NHCl2+ H2O→HOCl+ NH2Cl k4=2.3X X 103h
-1 NH3Cl++Br
- NH3Br
++Cl
- k12=1.8 X 108M
-1h
-1
NH2Cl + NH2Cl→NHCl2+NH3 ak5=pH dependent NH3Br
+NH2Br+H
+ pKa=6.4
NHCl2+NH3→NH2Cl + NH2Cl k6=2.16 X 108M
-2h
-1 HOCl+Br-HOBr+Cl
- k13=5.6 X 106M
-1h
-1
NH2Cl + NHCl2→N2+3H++3Cl
- k7=55.0M-1
h-1 HOBr+NH3NH2Br+H2O k14=2.7 X 10
11M
-1h
-1
NHCl2+H2O→bI+2HCl k8=4.0 X 10
5M
-1h
-1 HOBr+NH2ClNHBrCl k15=1.0 X 109M
-1h
-1
te.intermedia reactive :I .C 25 at ,hM800k and
,hM100.4k ,hM105.2k :]HCO[k]COH[k]H[kk :Note
b12
3HCO
1233CO2H
127
H33HCO323CO2HH5
a
(Source: Vikesland et al., 2001)
Disinfectant Residual Reactions in Distribution Systems 2.1.2
Both free chlorine and NH2Cl experience temporal and spatial degradation in distribution
systems in the bulk water and at the pipe wall due to chemical and biological reactions.
Therefore, both the bulk water and the pipe wall create a demand for the disinfectant residual.
The bulk water demand comes from dissolved organic matter, ammonia and nitrite, and some
dissolved metallic compounds (e.g. ferrous ions and manganese). NH2Cl may also spontaneously
decay in the absence of other reactants by auto-decomposition, which includes a complex series
of reactions and accounts for approximately half of the NH2Cl loss (Duirk et al., 2005). The
stoichiometry of NH2Cl auto-decomposition can be characterized by the generalized reaction
(Vikesland et al., 1996) of:
3NH2Cl N2 + NH3 +3Cl- + 3H
+ Equation 2-1
Auto-decomposition is impacted by pH and Cl2/N ratio. It is recommended that the pH remain
above 8.3 and the Cl2/N mass ratio be approximately 4:1 to maintain a chloramine residual in
distribution systems (Wilczak et al., 2003a).
The pipe wall demand of the disinfectant residual is caused by biofilms, deposits of
corrosion byproducts, and organic matter adsorbed onto metal oxide or carbonate scales (Lu et
12
al., 1999; Rossman et al., 2001; Vikesland and Valentine, 2002). The reactions of chlorine with
the scales on the inner pipe surface are regarded as the main reason for the loss of secondary
disinfectant residual within distribution networks (DiGiano and Zhang, 2005). Vikesland and
Valentine (2002) have reported that iron corrosion products can catalyze monochloramine
degradation, and the catalytic activities of the oxides can be ranked in the following order:
magnetite > goethite (α-FeO(OH)) > lepidocrocite (-FeO(OH)) hematite (Fe2O3) >
ferrihydrite. Lu et al. (1999) observed that at steady state biofilms chlorine demand increased
with increasing bioeliminable dissolved organic carbon (BDOC) of the water and the
surface/volume ratio of the pipes. However, Hallam et al. (2002) concluded in their research that
the biofilm could slow down chlorine consumption because it prevented chlorine reaching the
pipe surface. Regarding the impacts of organic matter, Rossman et al. (2001) reported that
although the rate of chlorine consumption in the pipe was much greater than in the bottle, the
increased formation of haloacetic acids and trihalomethanes (by 15%) was observed due to the
reactions of chlorine with organic precursor materials associated with deposits on the pipe wall.
Substantial research to date either in the field or in the laboratory has concluded that the
rates of disinfectant residual degradation in distribution systems increase with increasing
temperature, initial disinfectant concentration, total organic matter (TOC), the presence of
corrosion product and residence time (Brereton and Mavinic, 2002; Lu et al., 1999; Vikesland
and Valentine, 2002; Wable et al., 1991). Generally, NH2Cl is much less reactive with corrosion
products compared to free chlorine, and thus the demand exerted by the corrosion products
occurs more slowly than with chlorine (Valentine et al, 2000). However, in chloraminated
distribution systems, accelerated chloramine decay has been observed in systems with high
levels of nitrification (Wolfe et al., 1990; Woolschlager et al., 2001). It is primarily due to the
redox reaction between monochloramine and nitrite, and the reaction can be described as:
NH2Cl + NO2- + H2O NO3
- + NH3 + HCl Equation 2-2
Based on the stoichiometric relationship between nitrite and NH2Cl as shown in Equation 2-2,
each mg/L of NO2--N will consume as much as 5 mg/L as Cl2 of residual (Zhang and Edwards,
2009).
13
Kinetic Models for Disinfectant Residual Decay in Distribution Systems 2.1.3
Disinfectant residuals in distribution systems are consumed by soluble components
within the bulk liquid phase (bulk decay) and materials associated with the pipe wall such as the
pipe surface, corrosion products and the biofilm (wall decay). Therefore, the overall disinfectant
degradation in distribution systems can be expressed as:
Equation 2-3
Some researchers have postulated that chlorine decay models could be based on first-order or
parallel first-order kinetics for bulk reactions, and, either zero-order or first-order kinetics for
pipe wall reactions (Digiano and Zhang, 2005; Kiéné et al., 1998; Wable et al. 1991;
Vasconcelos et al., 1997).
For chlorine decay in the bulk flow, Kiéné et al. (1998) described the kinetic equation as:
CKdt
dCb
Bulk
Equation 2-4
where Kb =bulk decay constant; C =chlorine concentration in the bulk. The kinetic constants for
reactions are dependent on the water quality, specifically on temperature and organic content of
the water (Kiéné et al., 1998). Therefore, Kb can be described as:
T
b
b eTOCaK
Equation 2-5
where a =1.8X106L/mg·h; b =6050; T =temperature; TOC represents the organic content.
For disinfectant loss as a result of pipe wall consumption, a film resistance model with a
first-order rate equation has been proposed by Rossman et al. (1994) as:
Equation 2-6
WallBulkTotal dt
dC
dt
dC
dt
dC
W1W
Wall
CKV
S
dt
dC
14
where S = surface area; V = pipe volume; KW1 = wall reaction rate constant; CW = disinfectant
concentration at the pipe wall. According to this model (Rossman et al., 1994), mass transfer of
chlorine towards the pipe wall is proportional to the difference in chlorine concentration between
the bulk liquid and the pipe wall. If it is assumed that chlorine reacts at the pipe wall and there is
no accumulation of chlorine at the wall, the rate of mass transfer of chlorine would be equal to
the rate of chlorine decay at the pipe wall. Therefore, the flux of chlorine to the wall can be
written as:
Equation 2-7
where KF = mass transfer coefficient; C = disinfectant concentration in the bulk. Solving
Equation 2-7 for CW and substituting it into Equation 2-6 gives:
Equation 2-8
where KW = overall wall decay constant combining the mass transfer constant, KF, and the wall
reaction constant KW1.
Rossman et al. (1994) also provided a relationship between the mass transfer coefficient,
KF (Equation 2-7) and the dimensionless Sherwood number (Sh) as:
Equation 2-9
Equation 2-10
3. 0.0 8(
)
1+0.0 (
)
2 3 for Re 2,300 Equation 2-11
Equation 2-12
Equation 2-13
)CC(KCK WFW1W
CK)V
S(C
)KK(
KK)
V
S(
dt
dCW
F1W
F1W
Wall
p
hFd
DSK
2,300Refor ;Re023.0 333.083.0 ScSh
v
ud pRe
D
vSc
15
where Sh = Sherwood number; dp = pipe diameter; Re = Reynolds number; Sc = Schmidt
number; L= pipe length; D = molecular diffusivity of chlorine in water; u = flow velocity in pipe;
v = kinematic viscosity of water. Based on Equations 2-9 to 2-13, a positive relationship between
disinfectant decay rates and flow velocity would be expected. Flow velocity can influence
turbulence, radial diffusion and boundary layer thickness by affecting the Reynolds and
Sherwood numbers. This helps to explain why some studies have shown that the rates of free
chlorine and chloramine decay can increase with rising flow velocities in cast iron, unlined
ductile iron and even PVC pipes (Digiano and Zhang, 2005; Hallam et al., 2002; Mutoti et al.,
2007; Westbrook and Digiano, 2009). In addition to the mass transfer rate, some researchers
such as Mutoti et al. (2007) have attributed some of the increase in disinfectant decay at high
velocity to the increase in the release rate of corrosion products from the pipe surface.
Kb and KW as shown in Equations 2-4 and 2-8 will exhibit different values for the various
pipe materials due to the different pipe wall reaction kinetics. Table 2-3 summarizes some of
these coefficients for free chlorine and chloramine in contact with different pipe materials.
Table 2-3 Free chlorine and chloramine model coefficients
Material Parameter Free chlorine value Parameter Chloramine value
PVC Kb 0.048 h
-1 Kb 0.020 h
-1
KW 0.007 h-1
KW 0.007 h-1
Lined cast
iron
Kb 0.084 h-1
Kb 0.017 h-1
KW 0.007 h-1
KW 0.011 h-1
Unlined cast
iron
Kb 0.172 h-1
Kb 0.147 h-1
KW 0.063 h-1
KW 0.084 h-1
(Source: Taylor et al., 2005) Note: both bulk decay for chlorine and chloramine followed parallel first order kinetics,
and wall decay for both species followed first order kinetics.
Cast iron Kb 0.04 h
-1 First order
KW 36.5 mg/m2/h Zero order
Ductile iron Kb 0.04 h
-1 First order
KW 0.07-0.26 h-1
First order
(Source: Digiano and Zhang, 2005)
16
2.2 Disinfection Byproducts
Haloacetic Acids 2.2.1
Haloacetic acids (HAA) form by the chlorination of natural organic matter (NOM). Since
the 1970s, considerable efforts have been made to understand HAA formation mechanisms by
using NOM or well-defined model compound precursors (Kanokkantapong et al., 2006a;
Kanokkantapong et al., 2006b; Morris, 1975; Reckhow et al., 1990; Rook, 1977). Reaction
mechanisms involved in HAA formation generally include oxidation, substitution, addition, and
hydrolysis (Morris, 1975). Among the nine possible chlorinated and/or brominated HAA
compounds, dichloroacetic acid (DCAA) and trichloroacetic acid (TCAA) are the most
commonly reported HAA species. The dominance of these species is dependent on the chlorine
dose, as well as the water’s temperature, pH, and the specific ultraviolet absorbance (SUVA,
defined as the UV absorbance of a water sample at a given wavelength normalized for total
organic carbon concentration). Higher HAA concentrations are favoured at an increased chlorine
dose, temperature and SUVA and/or at low pH. Species that are generally low in concentration
in water include bromochloroacetic acid (BCAA), dibromoacetic acid (DBAA),
monochloroacetic acid (MCAA), monobromoacetic acid (MBAA), bromodichloracetic acid
(BDCAA), chlorodibromoacetic acid (CDBAA) and tribromoacetic acid (TBAA). Toxicological
properties of the different HAA species depend on the extent and type of halogen substitution
(bromine or chlorine). Guidelines for Canadian Drinking Water Quality established the
Maximum Acceptable Concentration (MAC) for HAA5 (MCAA, DCAA, TCAA, MBAA, and
DBAA) in drinking water at 80 µg/L (Health Health Canada, 2008). The Stage 2 D/DBPR
regulated the Maximum Contaminant Level (MCL) of HAA5 at 60 µg/L (USEPA, 2006).
N-nitrosodimethylamine 2.2.2
Although the application of chloramines has effectively reduced the formation of
halogenated DBPs, it may cause some unintended changes in water quality, including increased
formation of nitrosamines (Choi et al., 2002; Najm and Trussell, 2001). N-nitrosodimethylamine
(NDMA) is a suspected human carcinogen. Currently, federal drinking water guidelines for
nitrosamines are under development in Canada and the United States. Health Canada has
established the guideline for NDMA in drinking water at a Maximum Acceptable Concentration
17
(MAC) of 40 ng/L (Health Canada, 2011). The Ontario Ministry of the Environment, Canada,
has set a MAC for NDMA of 9 ng/L (MOE, 2002), and the California Department of Health
Services has set a notification level of 10 ng/L for each of NDMA, N-nitrosodiethylamine
(NDEA) and N-nitrosodi-n-propylamine (NDPA), respectively (CDHS, 2009). Mechanistic
studies of nitrosamines, especially NDMA, have been widely implemented with the
improvement of sensitive analytical techniques (Cheng et al., 2006; Mitch et al., 2003a). Several
reaction pathways have been investigated under controlled conditions as follows.
a. NDMA Formation from the Reaction of Monochloramine and Dimethylamine
Two steps are included in this mechanism: a nucleophilic substitution of dimethylamine
(DMA) by monochloramine (NH2Cl) to form an unsymmetrical dimethylhydrazine (UDMH) and
subsequent oxidation of UDMH intermediate to NDMA (Choi and Valentine, 2002; Mitch et al.,
2003b; Mitch et al., 2005; Mitch and Sedlak, 2002). The reactions are given in Equations (2-14)
and (2-15):
NH2Cl + (CH3)2NH → (CH3)2NNH2 + H+ +Cl
+ Equation 2-14
(CH3)2NNH2 + 2NH2Cl + H2O → (CH3)2NNO + 2NH3 + 2H++ 2Cl
- Equation 2-15
The rate of UDMH formation is slow and increases with pH (Yagil and Anbar, 1962),
whereas the UDMH oxidation occurs nearly instantaneously but with low yields (<1%) (Mitch
and Sedlak, 2002). The overall reaction rate is therefore controlled by the first step, and it
requires a long reaction time for NDMA formation. This may explain the increased NDMA
formation in distribution systems, especially in dead-end regions where long retention times are
expected.
Schreiber and Mitch (2007) proposed two other NDMA formation pathways involving
NHCl2: a relatively slow reaction of NHCl2 with amine precursors in the presence of dissolved
oxygen, and a fast reaction involving reactive breakpoint chlorination intermediates (HNO)
which was produced from the hydrolysis of NHCl2. Normally, NHCl2 can be minimized by
controlling pH (> 8.5) and the Cl2/N mass ratio (< 5:1) (Schreiber and Mitch, 2006). In addition,
UDMH
18
increasing the temperature accelerates the auto-decomposition of NH2Cl, thereby decreasing
NDMA formation (Mitch et al., 2003a).
b. NDMA Formation from Other Precursors
Humic substances in natural organic matter and tertiary amines with dimethylamine
functional groups have been suggested to serve as additional sources for nitrosamine precursors
(Chen and Valentine, 2006; Mitch et al., 2003b; Siddiqui and Atasi, 2001). Chen and Valentine
(2007) reported that hydrophilic fractions of natural organic matter appeared to form more
NDMA than hydrophobic fractions, and basic fractions tended to have a larger NDMA formation
potential than acid fractions. NDMA formation from tertiary amines is actually attributed to the
formation of secondary amines which contain dimethylamine functional groups through a
dealkylation process (Ellis and Soper, 1954). The primary amine (monomethylamine), the
quaternary amine (tetramethylamine), and amino acids or proteins do not form a significant
amount of NDMA after chloramination (Mitch and Sedlak, 2002). Strong-base anion exchange
resin and cationic polyelectrolytes (Epi-DMA and polyDADMAC products) used as flocculation
aids, which contain organic nitrogen, have also been observed to produce NDMA by reactions
with chloramine or hypochlorous acid (Gerecke and Sedlak, 2003; Najm and Trussell, 2001;
Westerhoff and Mash, 2002). Industrial products containing dimethylamine functional groups
including fungicides (e.g. tetramethylthiuram disulfide (thiram)), and drugs (e.g. ranitidine) have
also exhibited NDMA formation potential during disinfection of drinking water (Graham et al.,
1996; Schmidt et al., 2006).
Disinfection By-Products in the Distribution System 2.2.3
The concentration of HAA9 and nitrosamines varies both spatially and temporally within
distribution systems. Factors affecting DBP formation in distribution systems include the
concentration and chemical properties of the precursors, water temperature, pH, bromide ion
concentrations, disinfectant type, dose and residual, as well as contact time (Baribeau et al.,
2001; Baribeau et al., 2006; Cowman and Singer, 1996; Liang and Singer, 2003; Nguyen et al.,
2002; Obolensky and Frey, 2002; Rossman et al., 2001; Singer et al., 2002).
The HAA9 concentration in chlorinated systems either increases or remains
approximately constant as water age increases, while remaining relatively constant in
19
chloraminated distribution systems (Baribeau et al., 2006). In both systems, the observed
decrease in HAA9 concentration following long residence time has been attributed primarily to
biodegradation (Chen and Weisel, 1998; Rodriguez et al., 2004). Research has also been
performed to investigate the impacts of pipe metals (mainly iron and copper) and their corrosion
byproducts on the occurrence of HAA9 in distribution systems. In most cases, less trichloroacetic
acid (TCAA) was present than dichloroacetic acid (DCAA) due to the consumption of
disinfectant residuals by these pipe metals and their corrosion byproducts (Chen and Weisel,
1998; Hassan et al., 2006; Li et al., 2008; Rossman et al., 2001; Tuovinen et al., 1984). It is
noteworthy that high chlorine losses that can be observed in some distribution systems due to
reactions with corrosion products will not necessarily reduce DBP formation rates. The organic
material sorbed to the pipe wall may actually increase DBP levels (Hassan et al., 2006; Korshin
et al., 1997; Rossman et al., 2001).
Direct reactions between pipe materials and HAA9 were reported by Holzalski et al.
(2001). They have observed that trichloro-, tribromo- chlorodibromo-, and bromodichloro- acetic
acid readily reacted with Fe0 via sequential hydrogenolysis (replacement of a halogen by
hydrogen). The reduction of HAA9 by Fe0
may account for some losses of HAA9 observed in
full-scale distribution systems where iron is abundant. However, when the percentage of Fe0 on
the pipe surface is far less than 10%, as might be in cases of old cast iron pipes, the degradation
of DBP by Fe0 may not be important (Zhang et al., 2004).
Limited data is available concerning the formation and fate of nitrosamines in full-scale
distribution systems. While no particular trend in NDMA concentration with increasing water
age has been observed in either chlorinated or chloraminated distribution systems by some
researchers (Baribeau et al., 2006), others have reported that NDMA production increased with
the retention time in distribution systems (Charrois et al., 2007; Wilczak et al., 2003b). Other
nitrosamines, including N-nitrosomorpholine (NMOR), N-nitrosopyrollidine (NPYR), N-
nitrosopiperidine (NPIP) and N-nitrosodiphenylamine (NDPhA), have also been detected in
some distribution systems at extreme ends (Zhao et al., 2006). Thus, residence time tends to
significantly affect the observed concentrations of nitrosamines in distribution systems. Two
studies have investigated the reduction potential of granular iron and nickel-enhanced iron on
NDMA, and they reported that iron can transform NDMA to dimethylamine (DMA) and
ammonium via catalytic hydrogenation (Gui et al., 2000; Odziemkowski et al., 2000). The
20
presence of a small amount of nickel (0.25%) plated onto the iron can greatly enhance the
NDMA transformation rate. However, it is uncertain if the same reaction can take place in
distribution mains where iron pipes are widely present and corroded. Schreiber and Mitch (2006)
have reported that copper may catalyze NDMA formation. However, details concerning factors
that impact copper catalysis and the catalytic potential of copper solid corrosion products on
NDMA formation are not available. Therefore, the role of iron, copper and other pipe metals as
well as their corrosion products on the fate of NDMA and other nitrosamines in distribution
systems needs to be further investigated.
2.3 Corrosion
Corrosion is a process that deteriorates materials through chemical reactions with their
environment. In drinking water distribution systems, the materials may be a metal pipe or fitting,
the cement in a pipe lining or a polyvinyl chloride (PVC) pipe. In recent years, considerable
efforts have been made to understand the mechanisms of metal release from pipes and corrosion
scales, the rates at which they occur under typical disinfectant residual concentrations and
hydraulic conditions, and effective strategies to control corrosion and/or metal release (e.g.
phosphate-based corrosion inhibitor addition).
Metal Dissolution and Precipitation in the Distribution System 2.3.1
Corrosion of metallic materials in distribution systems is an electrochemical process in
which the metallic material is oxidized to a cation at an anodic site according to the reaction of:
M Mn+
+ ne- Equation 2-16
The released electrons travel through the conducting electrolyte (e.g. water) to a site which acts
as the cathode. At the cathode, the most common reaction in distribution systems is the
acceptance of electrons by dissolved oxygen or by a disinfectant residual that is in contact with
the metal. The metal ions are then either released into the drinking water as corrosion products or
they may react with components present in the water (such as OH-, CO3
2-, Cl
-, and SO4
2-) to form
hydroxyl, carbonate, chloride, and sulfate complexes as well as solid phases that may precipitate
on the pipe surface. Table 2-4 summarizes the principal equilibrium reactions of copper in
drinking water.
21
Table 2-4 Relevant equilibrium reactions for copper in carbonated-bearing water
Log K* (25 °C)
A) Direct Oxidation of Copper Metal in Chlorinated Water
HOCl+Cu(s) = CuO(s) + H+ + Cl
- 31.78
HOCl+Cu(s)+H2O= Cu(OH)2(s) + H+ + Cl
- 30.5
HOCl+Cu(s)+HCO3-= CuCO3(s) + H2O + Cl
- 38.41
2HOCl+2Cu(s)+HCO3-= Cu2(OH)2(CO3)(s) + H
+ + 2Cl
- 71.84
B) Dissolution of Solid Phases
CuCO3(s) = Cu2+
+ CO32-
-11.5
Cu2(OH)2(CO3)(s) = 2Cu2+
+ 2OH- + CO3
2- -33.3
Cu(OH)2(s) = Cu2+
+ 2OH- -19.32
CuO+H2O = Cu2+
+ 2OH- -20.35
C) Metal-Ligand Complexation
Cu2+
+ CO32-
= CuCO3(aq) 6.77
Cu2+
+ 2CO32-
= Cu(CO3)22-
(aq) 10.2
Cu2+
+ OH- = CuOH
+(aq) 6.5
Cu2+
+ 2OH- = Cu(OH)2(aq) 11.8
Cu2+
+ Cl- = CuCl
+ 0.2
(Source: Hong and MacAuley, 1998)
Normally, tenorite [CuO], cupric hydroxide [Cu(OH)2], and malachite [Cu2CO3(OH)2]
are dominant solid phase for copper. Iron and lead corrosions in distribution systems are also
complex processes which may involve many equilibrium reactions between dissolution and
precipitation, between soluble complexes and ion pairs, and even redox reactions. The common
species found in the iron pipe passive layers include goethite [α-FeOOH], lepidocrocite [γ-
FeOOH], magnetite [Fe3O4], ferric hydroxide [Fe(OH)3]. As for lead, PbCO3 and
Pb3(CO3)2(OH)2 are primary solid phases in drinking water. When free chlorine is present,
divalent lead can be further oxidized to tetravalent lead dioxide (PbO2) (Switzer et al., 2006).
The dissolution of the dissolved metal from the solid phase is controlled by the solubility of the
solid species, the concentration of ions between the pipe wall and bulk liquid, flow velocity, and
the water quality. The influence of some of these factors will be reviewed in Section 2.3.2.
Factors Affecting Metal Levels at the Tap 2.3.2
pH and alkalinity are the most significant factors influencing metal levels in distribution
systems. pH will influence the solubility of corrosion byproducts, thus affecting metal
22
concentrations at the tap. Normally, lead, copper and iron solubility decreases with elevated pH
(McNeill and Edwards, 2001; Sarin et al., 2003; Schock and Gardels, 1983; Schock, 1989; Xiao
et al., 2007). In the neutral pH range (i.e., pH from 6 to 8), the equilibrium concentration of lead
could vary by a factor of at least 5 to 10 per pH unit (Schock and Wagner, 1985). Solubility
models show that the lowest lead levels occur when pH is around 9.8 (Schock, 1989; Schock and
Gardels, 1983). Examination of copper levels under the USEPA Lead and Copper Rule from the
experience of 361 utilities revealed that the average 90th-percentile copper levels were highest in
waters with pH below 7. and that no utilities with pH above 7.8 exceeded the USEPA’s action
level for copper of 1.3 mg/L (Dodrill and Edwards, 1995). For iron, Sarin et al. (2003) have
reported that increasing pH from 7.6 to 9.5 can reduce the amount of iron released to water (<
0.25 mg/L).
The degree to which alkalinity affects the solubility of iron, copper and lead is different.
Lower iron corrosion rates and iron concentrations in distribution systems have been associated
with higher alkalinity due to the formation of the less soluble siderite species (FeCO3, Ksp =
2.110-11
] (Sarin et al., 2003). However, it has been demonstrated in many laboratories and in
utilities that the release of copper corrosion products is worse at higher alkalinity due to the
formation of soluble cupric bicarbonate and carbonate complexes (Edwards et al., 1996; Knox et
al., 2005; Schock et al., 1995). For lead, the impact of alkalinity is dependent on the form of lead
carbonate present on the pipe surface. In a study by Schock (1990), when PbCO3 (Ksp =7.4
10-14
) was present, increasing alkalinity above 50 mg/L CaCO3 reduced lead solubility in the pH
range of approximately 6 to 7.3. The author also found that when Pb3(CO3)2(OH)2 (Ksp =10-18.8
)
was present, lead solubility would not be significantly affected by alkalinity in this pH range, but
the solubility would decrease if the pH was increased above 7.3.
Limited research has been conducted on the role of disinfectant type on corrosion.
LeChevallier et al., 1990) reported that free chlorine was more corrosive to iron than NH2Cl in
deionized water with extremely low alkalinity and hardness. Cantor et al. (2003) also found an
increased iron corrosion rate in an iron pipe loop in the presence of free chlorine compared to
NH2Cl. Free chlorine residual was also shown to increase copper corrosion at low pH because of
the increased oxidizing strength of HOCl (Boulay and Edwards, 2001; Cantor et al., 2003;
Rahman et al., 2007). Limited information has been reported concerning the effect of chloramine
on copper. Rahman et al. (2007) reported that the application of NH2Cl could reduce the
23
dissolution of copper pipes from 1.26 mg/L in the presence of free chlorine to 0.48 mg/L.
However, ammonia released from chloramine decomposition can be aggressive to copper and
copper alloys by forming stable copper-ammonia complexes (Cu(NH3)22+
, Log K at 25 °C =
7.47) (Boyd et al., 2009; Schock and Lytle, 1995). The high oxidation potential of free chlorine
is sufficient for the formation of lead dioxide [PbO2] (Ksp =10-66
). In contrast, divalent lead
solids, primarily in the form of Pb3(CO3)2(OH)2 (hydrocerussite, Ksp =10-18.8
), are favored under
chloramination conditions (Vasquez et al., 2006). In the fall of 2003 in Washington DC, a high
level of Pb2+
(48000 ppb) was detected in their distribution systems following their shift from
free chlorine to NH2Cl in 2000 (Renner, 2004). It was theorized that free ammonia combined
with a high concentration of nitrate may have synergistically driven the lead corrosion by
interfering with the formation of the passive layer on its metal surface (Edwards and Dudi, 2004;
Uchida and Okuwaki, 1998; Zhang et al., 2009a). The direct reaction of metallic lead with nitrate
can be illustrated with the formula:
NO3- + Pb NO2
- + PbO Equation 2-17
Corrosion Control 2.3.3
Released metal ions and particles may affect aesthetic quality of delivered water and/or
pose potential risk to human health. Therefore, for metal pipes, the drinking water guidelines
have established the maximum acceptable concentration (MAC) of lead, based on its health
effects in children, at 0.010 mg/L. For iron and copper, based on aesthetic considerations, their
MACs are set at 0.3mg/L and 1.0 mg/L, respectively (Health Canada, 2009).
To control metal corrosion, the addition of phosphate-based chemicals has been used as
one effective strategy. Phosphate-based corrosion inhibitors can be dosed as either
orthophosphoric acid, combinations of orthophosphoric acid and zinc orthophosphate,
polyphosphates, or blends of orthophosphoric acid and polyphosphate. The effectiveness of
orthophosphate as a corrosion control measure relies on the development of impervious solid
films to act as a barrier scale between the corroding metal and the water (Demora and Harrison,
1984; Singley, 1994). However, polyphosphates have a different mode of action. Polyphosphates
are strong chelating agents which actually reduce scaling and have been shown to be effective in
sequestrating Fe2+ ions to treat “red water” (Facey and Smith, 199 ; Maddison et al., 2001;
24
Williams, 1990). Several authors have demonstrated that either orthophosphate or polyphosphate
can reduce iron levels and prevent iron corrosion (Facey and Smith, 1995; Lytle and Snoeyink,
2002; Maddison et al., 2001; Sarin et al., 2003; Williams, 1990).
Orthophosphate may reduce copper release in short term by blocking active sites on the
copper surface and thus protecting metallic copper against oxidation (Dartmann et al., 2004).
However, in long term, orthophosphate has been shown to interfere with the natural development
of malachite scales [Cu2(OH)2CO3], leading to an increased copper concentration in the water
(Cantor et al., 2000). As well, the performance of polyphosphate at a practical dosage (about 1
mg/L as P) on copper corrosion control can be unpredictable. It has been shown to either slightly
decrease or increase copper corrosion (Cantor et al., 2000; Colling et al., 1992; Edwards et al.,
2002).
Orthophosphate and polyphosphate have opposite effects on lead solubility. The
application of orthophosphate can decrease lead release through the formation of relatively
insoluble hydroxypyromorphite [Pb5(PO4)3OH] scales (Edwards et al., 1999). A 70% reduction
in lead release by orthophosphate has been reported (Edward and McNeill, 2002). In contrast,
polyphosphate might increase lead leaching into the water by forming lead polyphosphate
complexes (Leroy, 1993; MacQuarrie et al., 1997; Schock and Wagner, 1985).
2.4 Nitrification
It has been reported that nearly two thirds of chloraminated utilities in the United States
experience nitrification episodes (Wilczak et al., 1996). As a result of NH2Cl application as a
secondary disinfectant, ammonia is present in distribution systems and can be converted to
nitrate (NO3-) by ammonia-oxidizing bacteria (AOB) and nitrite-oxidizing bacteria (NOB)
through the formation of the nitrite (NO2-).
Nitrification in a distribution system can cause many potential water quality problems
(Wilczak et al., 1996; Zhang et al., 2009b), including:
the depletion of disinfectant residual;
DBP formation due to mitigation techniques;
the reduction in pH and alkalinity;
25
the formation of nitrite and nitrate.
and the increased autotrophic and heterotrophic microorganism growth.
Although the reduction in pH and alkalinity is not a potential threat to public health, it could
theoretically violate the USEPA Lead and Copper Rule (2008) and cause the dissolution of lead
and copper to unacceptably high levels if these designated water quality parameters fail to be
maintained. The simultaneous occurrence of nitrification and copper corrosion has been
observed within household plumbing systems (USEPA, 2002).
Factors affecting nitrification include pipe materials, phosphate, nutrients, pH,
disinfectant types and residuals, and temperature (Wilczak et al., 1996; Zhang and Edwards,
2009; Zhang et al., 2009a; Zhang and Edwards, 2010b; Zhang et al., 2008). Different pipe
materials can affect nitrification by providing a source of trace nutrients, acting as a toxic metals,
providing a support for attached growth and consuming disinfectant (Zhang et al., 2009b). A low
level of copper (1-10 ppb) from corrosion may facilitate nitrification, whereas an increase in
copper concentration (> 100 ppb) could inhibit nitrification (Zhang and Edwards, 2010b). Lead
materials are capable of producing nitrite and ammonia via electrochemical reactions with nitrate
(Edwards and Dudi, 2004; Uchida and Okuwaki, 1998; Zhang et al., 2009b), and the released
ammonia can support nitrification activity. Therefore, lead surfaces appear to be favored by
nitrifiers rather than other surfaces (Zhang et al., 2008). As such, increased lead corrosion rates
may stimulate nitrification by recycling ammonia from nitrate. Corroded iron surfaces are
favored by nitrifying bacteria because they can provide essential micronutrients for nitrifier
growth (Morton et al., 2005). Iron corrosion can also accelerate chloramine decay and release
ammonia to support the nitrifier population (Odell et al., 1996).
Water treatment practices to control nitrification include altering the chlorine-to-
ammonia dosing ratio and increasing the initial disinfectant dose. During chloramination, the
Cl2/N molar ratio is typically maintained between 0.6 to just below 1.0. Increasing the Cl2/N
ratio has been found to effectively mitigate the potential of nitrification by minimizing the
available ammonia to as close as possible to 1:1 (Lieu et al., 1993). Since free chlorine residual
is more effective in inactivating AOB than chloramine, periodic breakpoint chlorination is often
used to control nitrification in distribution systems (Odell et al., 1996; Wolfe et al., 1990).
Nitrifying bacteria have slow growth rates, and thus grow best with an extended detention time,
26
e.g. in large reservoirs or in dead-end regions of distribution systems. Therefore, operational
activities that shorten water age may be considered to control nitrification, including increased
daily turnover of water in reservoirs/storage tanks, installation of recirculation facilities, and the
use of small diameter pipes (Odell et al., 1996).
2.5 Summary of Research Gaps
The water industry’s move to consider changing the secondary disinfectant from free
chlorine to chloramine can have negative impacts on the delivered drinking water quality,
including unintended nitrosamine formation, nitrification and increased corrosion rates (e.g.
lead). However, the literature information regarding the behavior of HAA9 and nitrosamine
(especially NDMA) concentrations in distribution systems, particularly with respect to the
influence of pipe wall materials, the presence of corrosion inhibitors, flow conditions, and, to
some extent, biofilm formation, is very limited and inconsistent. Therefore, further studies are
needed to understand the fate (formation and/or degradation) of these DBPs in the complex
physiochemical and biological environment of distribution systems. Several questions that could
be asked include:
Are there differences in disinfectant degradation and HAA/NDMA formation for
different pipe materials treated with different phosphate-based corrosion
inhibitors?
Do metal age and water quality affect the efficacy of corrosion inhibitors and
disinfectant decay as well as HAA/NDMA formation?
How do flow conditions potentially affect secondary disinfectant stability and
NDMA formation?
What are the impacts of solid corrosion products on the fate of HAA/NDMA in
distribution systems?
Therefore, the current research was implemented to at least partially answer these questions and
bridge the aforementioned gaps thus aiding the water treatment industry to develop strategies
that minimize water quality degradation.
27
2.6 References
Baribeau, H., Prevost, M., Desjardins, R., and Lafrance, P. (2001) Changes in chlorine and DOX
concentrations in distribution systems. Journal American Water Works Association,
93(12), 102-114.
Baribeau, H., Boulos, L., Hileselassie, H., Crozes, G., Singer, P. C., Nichols, C., Schleisinger, S.
A., Gullick, R. W., Williams, S. L., Williams, R. L., Fountleroy, L., Andrews, S. A., and
Moffat, E. (2006) Formation and decay of disinfection byproducts in the distribution
system. Water Research Foundation and US EPA, Project # 2770, Denver, USA.
Boyd, G. R., Dewis, K. M., Korshin, G. V., Reiber, S. H., Schock, M. R., Sandvig, A. M., and
Giani, R. (2009) Effects of Changing Disinfectants on Lead and Copper Release (vol 100,
pg 75, 2008). Journal American Water Works Association, 101(2), 10-10.
Boulay, N., and Edwards, M. (2001) Role of temperature, chlorine, and organic matter in copper
corrosion by-product release in soft water. Water Research, 35(3), 683-690.
Brereton, J.A. and Mavinic, D.S. (2002) Field and material-specific simulated distribution
system testing as aids to understanding trihalomethane formation in distribution systems.
Canadian Journal of Civil Engineering, 29(1), 17-26.
California Department of Public Health (CDPH) (2009). California Drinking Water: NDMA and
Other Nitrosamines - Drinking Water Issues.
http://www.cdph.ca.gov/certlic/drinkingwater/Pages/NDMA.aspx (accessed September
15, 2011).
Cantor, A. F., Denig-Chakroff, D., Vela, R. R., Oleinik, M. G., and Lynch, D. L. (2000) Use of
polyphosphate in corrosion control. Journal American Water Works Association, 92(2),
95-102.
Cantor, A.F., Park, J.K. and Vaiyavatjamai, P. (2003) Effect of chlorine on corrosion in drinking
water systems. Journal American Water Works Association, 95(5), 112-123.
28
Charrois, J. W. A., Boyd, J. M., Froese, K. L., and Hrudey, S. E. (2007) Occurrence of N-
nitrosamines in Alberta public drinking-water distribution systems. Journal of
Environmental Engineering and Science, 6(1), 103-114.
Chen, Z. and Valentine, R.L. (2006) Modeling the formation of N-nitrosodimethylamine
(NDMA) from the reaction of natural organic matter (NOM) with monochloramine.
Environmental Science & Technology, 40(23), 7290-7297.
Chen, Z. and Valentine, R.L. (2007) Formation of N-nitrosodimethylamine (NDMA) from humic
substances in natural water. Environmental Science & Technology, 41(17), 6059-6065.
Chen, W.J. and Weisel, C.P. (1998) Halogenated DBP concentrations in a distribution system.
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40
Free Chlorine Degradation and HAA Formation in Two 3
Water Matrices in Contact with Three Metal Materials
Experiments that formed the basis for Chapters 3 and 4 compared the kinetics of metal
corrosion, chlorine degradation and HAA formation for iron, copper and lead treated with
different corrosion inhibitors on a “survey” basis. Metal release kinetics data will be discussed in
Chapter 4 due to the quantity of data. Because most of the results from this survey study
regarding each metal material can be explained well by previously proposed theory and/or
mechanisms in the literature, no further mechanistic investigation was explored in these two
chapters. However, efforts to compare the impacts of corrosion inhibitors on free chlorine decay
and HAA formation among three metal materials have not been reported previously, and thus
this information is relatively new to the field. The observations concerning enhanced HAA
formation in the presence of copper also laid the groundwork for further investigations into
copper catalysis as reported in Chapters 5 and 6.
The results of this chapter in Sections 3.3.1.1, 3.3.2.1, 3.3.2.3, and 3.3.3.1 that deal with
chlorine degradation and HAA formation for fresh metal coupons in one type of water
(Mannheim Water) have been published as part of:
Zhang, H., Andrews, S. A. (2012) Effects of Phosphate-based Corrosion Inhibitors on the
Kinetics of Chlorine Degradation and HAA Formation in Contact with Three Metal
Materials. Canadian Journal of Civil Engineering, 39, 44-54.
Data from similar experiments using fresh metal coupons to investigate free chlorine decay and
HAA formation in another water matrix (Britannia Water) and experiments using pre-corroded
coupons for both tested water matrices are included in this chapter for completeness.
Results from this chapter focus on the research gaps “Are there differences in disinfectant
degradation and HAA formation for different pipe materials treated with different phosphate-
based corrosion inhibitors?” and “Do metal age and water quality affect disinfectant stability and
HAA formation?”
41
Abstract
Pipe materials and their corrosion products can significantly affect disinfectant residual
stability. However, with increasing application of corrosion inhibitors in distribution systems for
corrosion control, no relevant information has been found in literature about the impacts of
corrosion inhibitors and/or their interactions with metal surfaces on disinfectant degradation and
disinfection byproduct formation for both newly installed and aged pipe systems. Therefore, this
chapter investigated and compared free chlorine degradation and HAA9 formation in the absence
and presence of orthophosphate and polyphosphate for two water matrices containing fresh or
aged ductile iron, lead and copper coupons. Material-specific formation potential (MS-FP) and
material-specific simulated distribution system (MS-SDS) tests were applied at bench scale.
Regardless of metal age and water types, the reactivity of metal materials with free chlorine
followed the sequence of Fe> Cu> Pb. The addition of phosphate-based corrosion inhibitors
generally increased HOCl degradation for fresh iron coupons, but decreased HOCl decay only
for fresh copper coupons. Generally, these phosphate-based corrosion inhibitors did not impact
HAA formation. For fresh copper coupons in both investigated water matrices, HAA formation
was enhanced in the presence of high levels of copper ions, indicating possible catalytic potential
of copper on HAA formation. In addition, DCAA was the dominant species observed for water
in contact with fresh metal coupons, whereas either MCAA or DCAA formation was favored
over other species in water containing corroded metal coupons.
Keywords: Free chlorine; Haloacetic acids; Iron; Copper; Lead; Corrosion inhibitor
42
3.1 Introduction
To maintain microbial water quality in distribution systems, secondary disinfectants are
applied. Free chlorine is currently the most widely used secondary disinfectant in the distribution
systems due to its low cost and effectiveness. In distribution systems, free chlorine may
experience temporal and spatial degradation due to chemical and biological consumption that
occurs in the bulk water and on the pipe wall. At least 0.2 mg/L of free chlorine residual should
be maintained to reduce the possible occurrence of further contamination (USEPA, 1989).
Factors affecting disinfectant residual stability in distribution systems include temperature,
concentration of organic matter, biofilms, hydraulics, and infrastructure such as pipe materials,
corrosion products, and pipe service age (Lu et al., 1999; Rossman et al., 2001; Brereton and
Mavinic, 2002; Haas et al., 2002; Hallam et al., 2002; Vikesland and Valentine, 2002; Al-Jasser,
2007).
Haloacetic acids (HAA) are formed by the chlorination of organic materials.
Dichloroacetic acid (DCAA) and trichloroacetic acid (TCAA) are the most commonly-reported
HAA species. Other species, generally at lower levels in low-bromide water, include
bromochloroacetic acid (BCAA), dibromoacetic acid (DBAA), monochloroacetic acid (MCAA),
monobromoacetic acid (MBAA), bromodichloracetic acid (BDCAA), chlorodibromoacetic acid
(CDBAA) and tribromoacetic acid (TBAA). The toxicological properties of HAAs depend on
the extent and type of halogen substitution (bromine or chlorine). The Stage 2 Disinfectants and
Disinfection Byproducts Rules D/DBPR regulated the Maximum Contaminant Level (MCL) of
HAA5 (MCAA, DCAA, TCAA, MBAA, and DBAA) at 0 μg L (USEPA, 2006). Guidelines for
Canadian Drinking Water Quality established the Maximum Acceptable Concentration (MAC)
for HAA5 in drinking water at 80 μg L (Health Health Canada, 2008).
Precursors of HAAs are primarily natural organic matter, including humic and fulvic
substances (Kanokkantapong et al., 2006). During chlorination, the formation of DCAA and
TCAA is significant, and the dominance of both species is dependent on chlorine dose,
temperature, pH, and specific ultraviolet absorbance (SUVA) of water. More halogenated species
are favored at increasing chlorine dose, temperature and/or SUVA or at lower pH. In distribution
systems, the concentrations of HAA9 vary both spatially and temporally. HAA9 concentrations in
chlorinated systems generally increase or remain approximately constant as water age increases,
43
while remaining relatively constant in chloraminated distribution systems (Baribeau et al., 2006).
In both systems, observed decreases in HAA9 concentration following long residence times have
been attributed primarily to biodegradation (Chen and Weisel, 1998; Rodriguez et al., 2004).
It is widely accepted that disinfectant residuals in distribution systems increase corrosion
rates, as secondary disinfectants serve as potential oxidants for distribution system pipes and
related plumbing materials. In recent years, considerable efforts have been made to understand
the mechanisms of metal release from pipes and corrosion scales, the rates at which they occur
under the influence of disinfectant residuals and hydraulic conditions. Effective strategies to
control corrosion and/or metal release have also been developed. One such strategy is applying
phosphate-based corrosion inhibitors in the form of either orthophosphoric acid, combinations of
orthophosphoric acid and zinc orthophosphate, polyphosphates, or blends of orthophosphate and
polyphosphate. Generally, orthophosphate forms an impervious solid film as a barrier between
the corroding metal and the water (Demora and Harrison, 1984; Singley, 1994), whereas
polyphosphates are strong chelating agents which have been successfully used for reducing
scaling and sequestering Fe2+ ions to treat “red water” (Facey and Smith, 199 ; Maddison and
Gagnon, 1999; Williams, 1990).
It can be hypothesized that the application of phosphate-based corrosion inhibitors which
significantly affect corrosion rates may, in turn, impact disinfectant degradation rates and the
formation kinetics of DBPs. The purpose of this study was to investigate how pipe materials
(ductile iron, copper and lead) and phosphate-based corrosion inhibitors (orthophosphate and
polyphosphate) influence free chlorine degradation and HAA formation. These pipe materials are
widely present in either water mains (such as ductile iron) or household plumbing (such as
copper), and lead may be leached into drinking water from old lead service lines, soldered joints
and brass plumbing fittings (Health Canada, 2009). Given that no reports have been found that
compare the behavior of free chlorine degradation and HAA formation for all three metal
materials, understanding the stability of free chlorine residual and the fate of HAA compounds
under the influence of these pipe materials and their phosphate-based corrosion preventive
strategies will benefit utilities and households to develop effective strategies to control the
deterioration of water quality in both distribution system water mains and premise plumbing.
44
3.2 Materials and Methods
Reagents and Materials 3.2.1
All chemicals used in this study were ACS grade or higher. The chlorine dosing solution
(approximately 3500 mg/L as Cl2) was prepared by diluting a concentrated solution of sodium
hypochlorite (NaOCl, 6%, VWR) in Milli-Q water. The target disinfectant concentration applied
in the tests was achieved by spiking free chlorine dosing solution into 1 L test water.
Orthophosphate (Na3PO4) and polyphosphate ([Na(PO2)]6) were selected to investigate their
impacts on metal corrosion in this study. They were prepared at a concentration of 500 mg/L as P
and were kept in the dark at 4ºC. The targeted dosages of these corrosion inhibitors were 1 mg/L
as P.
Test coupons of ductile iron, copper and lead were purchased from Metal Samples Co.,
Alabama, US. The size of these coupons is 1 2”3”1 1 ”. For experiments with fresh coupons,
before each batch of experiments, any corrosion products were removed from the coupons by
polishing with 60-grit sandpaper followed by 120-grit sandpaper, and then rinsing with deionized
water and acetone followed by Milli-Q water (Scholze et al., 1994). Polished coupons were
submerged into the test water to simulate contact with new pipes without any impacts from
service age. In addition, to investigate the aged pipe environment, some coupons of each pipe
material were pre-conditioned in tap water in the absence and presence of orthophosphate. The
initial free chlorine concentration for conditioning purposes was 10 mg/L. After each 24 hour
reaction period for the iron coupons and each 48 hour reaction period for the copper and lead
coupons, the free chlorine residual was measured and then the water was replaced with fresh tap
water dosed with 10 mg/L free chlorine. The conditioning process proceeded until the 24 hour
chlorine residual for each combination of metal type and corrosion inhibitors was approximately
constant (< 4.0 mg/L chlorine decay for the iron coupons alone, 6.8 mg/L for the iron coupons
treated with orthophosphate, 2.7 mg/L for all of the copper coupons, and 1.9 mg/L for all of the
lead coupons). The measured 24 hour chlorine residuals for iron, copper and lead are shown in
Appendices Figure 10-14. Orthophosphate was selected since it effectively controlled metal
release for fresh metal coupons in short term, especially for copper and lead.
45
Unchlorinated post-filtration water was collected from Mannheim Water Treatment Plant
(MWTP) and Britannia Water Purification Plant (BWPP), Ontario. Water quality parameters are
listed in Table 3-1.
Table 3-1 Summary of water quality parameters for post-filtration water
Parameters
Values
Mannheim
Water
Britannia
Water
pH 7.5 ± 0.2 6.6±0.5
Alkalinity (mg/L) 187±33 12±2
UV254 (cm-1
) 0.058±0.008 0.065±0.005
TOC (mg/L) 4.5±0.3 3.3±0.2
SUVA (L/mg·cm-1) 0.013±0.002 0.019±0.0002
Bromide (μg L) 65.0±15.5 15.8±0.2
Chloride (mg/L) 84.5±2.5 3.1±0.6
Sulfate (mg/L) 35.0±2.0 23.8±0.8
Cl-:SO4
2- ratio 2.4±0.1 0.13±0.04
Note: TOC, total organic carbon; SUVA, specific ultraviolet absorbance.
Experimental Procedures 3.2.2
“Material-specific” simulated distribution system (MS-SDS) and material-specific
formation potential (MS-FP) tests were applied in this study. Details of MS-SDS procedures
have been described by Brereton and Mavinic (2002). They consist of incubating metal coupons
in water samples under conditions representative of actual field conditions in terms of reaction
time, pH, temperature, and disinfectant application. In MS-SDS tests, not only is disinfectant
interaction with precursors in the bulk water considered, but the influences of distribution system
pipe walls on disinfectant degradation and HAA generation are also examined. The MS-SDS
tests were performed for experiments employing corroded metal coupons, using an initial free
chlorine concentration of approximately 5.5 mg/L to ensure detectable disinfectant residuals (0.2
mg/L) after 24 hours. Experiments with fresh metal coupons employed material-specific
formation potential (MS-FP) tests that were performed similarly to the MS-SDS tests except that
higher concentrations of free chlorine (12.3 mg/L) were applied than would normally be
encountered during typical water treatment to meet the high chlorine demand of fresh metal
coupons and their corrosion products and ensure detectable chlorine residuals after 24 hours as
well.
46
All of the MS-SDS and MS-FP experiments were performed in 1 L amber bottles with
PTFE lined caps and at room temperature (21±2 ºC). The reaction bottles were made chlorine
demand free before use by soaking them in a concentrated sodium hypochlorite solution (~1000
mg/L as Cl2) for at least 24 hours. Thereafter, the bottles were rinsed thoroughly with deionized
water and distilled water, and were heated at 250 °C for at least 4 hours. Metal coupons were
suspended in the bottles using nylon threads. To maximize the contact of water with coupon
surface, a Big Bill Orbital shaker was used (Barnstead International, USA) to maintain a gentle
mixing at a speed of 25 rpm. After each designated reaction time, 50 mL of water were
withdrawn from the test bottles for the analysis of disinfectant residual, HAA and metal
concentrations. Seven to ten time points were employed for each of the kinetic studies. As will
be discussed in Section 4.3.1 (Chapter 4), metal concentrations reached equilibrium after 10
hours, and the reduction of water volume due to sequential sampling did not significantly affect
equilibrated metal concentrations. Separate control tests were also performed to confirm that
there were also no losses of HAA compounds due to volatility during the sequential sampling
events.
Two sets of kinetic experiments were performed to test the reproducibility of the results,
and duplicate tests were conducted in each set of experiments. All of the tests also included
control samples, which were prepared the same as test samples but without metal coupons.
Analytical methods that were employed to examine water quality and analytes of interest are
summarized in Table 3-2.
Table 3-2 Summary of analytical methods
Analyte Unit Instrument /procedure Reference method
TOC mg/L O1-Analytical TOC analyzer SM 5310 C
pH pH meter
UVA254 cm-1
Hewlett Packard 8452A Diode Array UV
spectrophotometer SM 5910B
Alkalinity mg/L Titration SM2320B
Chlorine mg/L Hach DR2700 Spectrophotometer SM 4500-Cl G
Anions µg/L Dionex DX-300 Series Ion
Chromatography System SM 4110 B
Haloacetic acids µg/L Hewlett Packard 5890 Series II Plus Gas
Chromatography SM 6251
Metal (Pb, Cu, and Cu) mg/L Varian SpectrAA.20 SM 3111B
Note: TOC, total organic carbon; SM represents Standard Methods for the Examination of Water and Wastewater
(APHA, AWWA, WEF, 2005).
47
3.3 Results and Discussion
The concentrations of free chlorine, released metal and HAA9 were monitored by
periodically withdrawing and analyzing samples from the test bottles. Using these data, it was
possible to compare free chlorine decay rates and HAA formation among the different
combinations of metals and corrosion inhibitors.
Free Chlorine Degradation 3.3.1
3.3.1.1 Fresh Coupons in Two Water Matrices
Figure 3-1 illustrates the observed free chlorine decay as a function of time for iron,
copper and lead coupons in the presence and absence of corrosion inhibitors in Mannheim
Water. Free chlorine degraded more quickly when in contact with all three pipe materials than
for the control samples without metal coupons, regardless of the type of corrosion inhibitors
present, although the effect was small in the presence of lead. Compared with the presence of
other metals, free chlorine consumption in the presence of iron coupons was much more rapid,
and almost no chlorine residual detected after 24 hours (Figure 3-1a). Therefore, to ensure
detectable chlorine residual at 24 hours (0.2 mg/L), an initial chlorine concentration of 12.3
mg/L was applied in all of the MS-FP bottles, even though it would be higher than observed
during normal applications.
The effects of the various experimental conditions on chlorine decay for copper coupons
can been seen in Figure 3-1b. The presence of corrosion inhibitors significantly reduced free
chlorine degradation rate in comparison with copper coupons in the absence of corrosion
inhibitors. Two possibilities may account for this phenomenon. One was due to rapid
consumption of free chlorine by copper corrosion itself. Without the protection from corrosion
inhibitors, the copper surface was more readily available to react with chlorine. The other
possibility was that dissolved copper ions could either directly enhance free chlorine degradation
to produce oxygen and chloride as final degradation products (OCl- O2+Cl
-; Gray et al., 1977)
or increase the reactivity of DBP precursors (Blatchley et al., 2003; Fu et al., 2009). Both of
these options would lead to faster consumption of free chlorine in the presence of pure copper
alone. No obvious impacts of corrosion inhibitors on free chlorine decay were observed for lead
coupons (Figure 3-1c). They appeared to have similar free chlorine decay behavior in the
48
presence of these corrosion inhibitors compared with metal coupons alone. It should be noted
that error bars, which represent the measured maximum and minimum values, cannot be seen in
Figure 3-1 since they were small (0.0 mg/L ~ to 0.2 mg/L) relative to the measured chlorine
residuals.
Figure 3-1 Free chlorine decay for fresh metal coupons with time in the presence and absence of
corrosion inhibitors in one set of experiments, n=2; (a) iron, (b) copper, (c) lead.
The rates of free chlorine decay were obtained by fitting a pseudo first-order decay
equation to the chlorine data from each experiment. In Ct = C0 exp (-kt), Ct is the free chlorine
concentration (mg/L) at time t, C0 the free chlorine concentration (mg/L) at time zero, and k the
0
3
6
9
12
15
0 10 20 30
Fre
e c
hlo
rin
e (
mg
/L)
Fe
Fe+polyphosphate
Fe+orthophosphate
Control
0
3
6
9
12
15
0 20 40 60 80 100 120
Fre
e c
hlo
rin
e (
mg
/L)
Cu
Cu+polyphosphate
Cu+orthophosphate
Control
b
c
a
0
3
6
9
12
15
0 20 40 60 80 100 120
Fre
e c
hlo
rin
e (
mg
/L)
Time (hours)
Pb
Pb+polyphosphate
Pb+orthophosphate
Control
49
overall first-order decay constant (h-1
) which is the sum of a bulk constant and a wall constant: k
= kb + kw. In this equation for k, kb is the bulk first-order chlorine decay constant (h-1
) and kw the
wall first-order chlorine decay constant (h-1
). The bulk water reaction in the MS-PF tests is the
same as would occur in the control bottles containing no metal coupons, thus the wall decay
constant is the difference between the overall decay constant and the bulk decay constant. Table
3-3 summarizes chlorine decay constants, including overall decay constants (k) and wall decay
constants (kw) for all three metal materials and the different corrosion inhibitors that were tested.
The variability in the decay constants as shown in Table 3-3 were likely attributed to the
different coupon surface conditions between sets of experiments even though all of the corroded
coupons were polished following the same procedures described in Section 3.2.2 before each
new set of experiments.
Table 3-3 Chlorine decay constants for fresh metal coupons in the presence and absence of
corrosion inhibitors with Mannheim Water (n=4)
Material Corrosion inhibitors k (h
-1) kw (h
-1)
Average Stdev Average Stdev
Iron
None 0.1451 0.0006 0.1402 0.0014
Polyphosphate 0.1834 0.0115 0.1789 0.0121
Orthophosphate 0.1668 0.0016 0.1620 0.0025
Copper
None 0.0292 0.0064 0.0239 0.0051
Polyphosphate 0.0169 0.0071 0.0117 0.0056
Orthophosphate 0.0137 0.0052 0.0084 0.0037
Lead
None 0.0078 0.0005 0.0029 0.0013
Polyphosphate 0.0075 0.0006 0.0030 0.0012
Orthophosphate 0.0078 0.0011 0.0029 0.0020
Control
(water)
None 0.0049 0.0008
NA Polyphosphate 0.0045 0.0006
Orthophosphate 0.0049 0.0009
Note: NA-not available.
In order to evaluate the interactive effects of corrosion inhibitors and metal surface type
on chlorine decay in Mannheim Water and to determine whether the treatment factor (the
different corrosion inhibitors) had a significant influence on the response factor (free chlorine
wall decay constants), a single-factor ANOVA test at a confidence level of 95% was applied to
the wall decay constants that were determined for each metal material. When the ANOVA test
signified statistically significant impacts of corrosion inhibitors on free chlorine wall decay, a
Fisher’s Least Significant Difference (LSD) test was applied to further determine if significant
50
differences existed between each pair of treatments at a 95% of confidence level (Montgomery,
2000).
As shown in Table 3-3, the reactivity of three metal materials with free chlorine followed
the sequence in decreasing order: iron, copper, and lead. Iron had wall decay constants two
orders of magnitude higher than for the bulk decay, irrespective of the type of corrosion
inhibitors present, so chlorine decay was predominant on the pipe wall. The ANOVA test
determined that there were significant differences in free chlorine decay among the three
treatments (p-value of 0.02), and the results of the LSD tests indicated that the corrosion
inhibitors significantly increased the free chlorine wall decay.
As was observed with iron, copper consumed free chlorine primarily by wall reactions,
and corrosion inhibitors exhibited significant impacts on free chlorine degradation (p-value of
0.01). However, the extent of the contribution of wall reactions to the overall chlorine decay was
dependent on the type of corrosion inhibitors present. The relative sequence of impact was: no
corrosion inhibitor < polyphosphate < orthophosphate. When LSD tests were used to compare
free chlorine wall decay constants in the presence of each type of corrosion inhibitors with those
for copper coupons alone, the results indicated that corrosion inhibitors significantly decreased
free chlorine degradation. This may have been from the formation of protective scales and/or
decreasing released copper concentrations, thereby making the copper surface less available to
react with free chlorine and/or decreasing copper catalysis.
In contrast to the results for iron and copper, chlorine consumption was dominated by
bulk water decay for lead coupons, for which their bulk water decay coefficients were at least
40% higher than their wall decay constants. Corrosion inhibitors did not exhibit significant
impacts on free chlorine wall decay relative to lead coupons in the absence of corrosion
inhibitors (p-value >0.05).
The relative reactivity of three metal materials with HOCl in Britannia Water had a
similar sequence to that observed in Mannheim Water. Namely, fresh iron coupons were most
reactive with free chlorine among three metal materials followed by copper, and then lead. The
addition of phosphate-based corrosion inhibitors significantly increased chlorine decay rates for
iron coupons, but reduced chlorine decay for copper and lead coupons.
51
The impacts of water quality and corrosion inhibitors on chlorine decay in the two water
matrices were determined using a two-factor ANOVA test at a confidence level of 95%. The
results of the ANOVA test (p-values) are shown in Table 3-4, along with the pseudo-first-order
HOCl decay constants.
Table 3-4 Comparison of HOCl overall decay constants (h-1
) for fresh coupons between two
water matrices (n=4)
Material Corrosion
inhibitors
Mannheim
Water
Britannia
Water
p-value (=5%)
Water
quality
Corrosion
inhibitors
Iron
None 0.1451±0.0006 0.1555±0.0305
0.02 0.001 Polyphosphate 0.1834±0.0115 0.2446±0.0040
Orthophosphate 0.1668±0.0016 0.1671±0.0041
Copper
None 0.0292±0.0064 0.0257±0.0075
0.23 0.01 Polyphosphate 0.0169±0.0071 0.0123±0.0011
Orthophosphate 0.0137±0.0052 0.0093±0.0009
Lead
None 0.0078±0.0005 0.0083±0.0021
0.25 0.27 Polyphosphate 0.0075±0.0006 0.0050±0.0006
Orthophosphate 0.0078±0.0011 0.0066±0.0023
Control
(water)
None 0.0049±0.0008 0.0038±0.0003
0.04 0.52 Polyphosphate 0.0045±0.0006 0.0034±0.0001
Orthophosphate 0.0049±0.0009 0.0041±0.0007
Initial free chlorine concentration: 12.3 mg/L as Cl2
The ANOVA test results suggest that in Britannia Water chlorine decay constants for
fresh iron coupons significantly increased (p-value 0.02) relative to Mannheim Water, while the
bulk water control samples in Britannia Water had slightly decreased chlorine decay constants
(p-value 0.04). Unexpectedly, fresh copper and lead coupons had statistically the same chlorine
degradation rates in each of the two water matrices. Since Britannia Water had lower pH and
alkalinity than Mannheim Water (Table 3-1), Britannia Water would be more corrosive to metal
coupons than Mannheim Water, and thus more rapid free chlorine degradation should have been
observed in Britannia Water than in Mannheim Water. However, this was only observed for
fresh iron coupons and the differences were small. Differences in the chlorine bulk decay in the
two water matrices can be likely attributed to their different TOC levels (Hallam et al., 2003;
Powell et al., 2000). In general, the higher the TOC concentration, the more rapidly free chlorine
will degrade. Therefore, it is consistent that the free chlorine bulk decay rates in Mannheim
Water were statistically higher than those in Britannia Water, regardless of the presence and the
types of corrosion inhibitors.
52
The ANOVA test results in Table 3-4 also indicate that in both water matrices the
presence of corrosion inhibitors significantly increased free chorine decay for fresh iron coupons
(p-value 0.001), while for copper coupons these corrosion inhibitors statistically decelerated
chlorine degradation (p-value 0.01). However, free chlorine degradation did not show any
statistical difference in the absence and presence of these corrosion inhibitors (p-value 0.27) for
solutions with fresh lead coupons and the bulk water control samples. In general, the results of
this two-factor ANOVA test provide new information concerning the effects of two different
water quality and corrosion inhibitors on free chlorine degradation.
3.3.1.2 Corroded Coupons in Two Water Matrices
Figure 3-2 compares the pseudo-first order free chlorine decay constants for three types
of corroded metal coupons in contact with Mannheim Water. The reactivity of these materials
with free chlorine followed the general sequence of iron >copper >lead, although the decay rates
were similar for solutions with corroded copper and lead when there was no corrosion inhibitor
present. The overall decay rates of chlorine in solutions with corroded iron treated with
orthophosphate were 0.21±0.001 h-1
, and were not statistically different from the decay rates in
the absence of orthophosphate (0.15±0.067 h-1
).
Figure 3-2 Free chlorine overall decay constants for corroded coupons with Mannheim Water.
Initial free chlorine concentration 12.3 mg/L, error bars indicate standard deviation (n=4)
For corroded copper coupons alone, the overall decay constant was only 70% higher than
those in the bulk water control, indicating that bulk water reaction was dominant for free chlorine
consumption. In contrast, corroded copper coupons treated with orthophosphate had an overall
free chlorine decay rate 1.8 times higher than the bulk water control, suggesting that the metal
0.00
0.05
0.10
0.15
0.20
0.25
No inhibitor Orthophosphate
1st o
rder
overa
ll decay c
onsta
nt
(h
-1)
Fe Cu Pb Control
53
wall reactions dominated free chlorine degradation. Some studies have revealed that
orthophosphate may interfere with the aging processes of passivating films of copper, e.g., from
soluble Cu(OH)2 to malachite over the long term (Li et al., 2004), thereby releasing copper from
the aged pipes. An increased copper corrosion rate in the presence of orthophosphate was also
observed in this water matrix (from 0.4 mg/L in the absence of orthophosphate to 0.5 mg/L after
the addition of orthophosphate at 24 hours). In addition, copper has been reported to catalyze
free chlorine degradation (Fu et al., 2009b; Gray et al., 1977). Further experiments have also
been conducted in this research to investigate factors affecting copper catalysis on chlorine
degradation (Chapter 5) and the extent of copper catalysis was shown to increase with increasing
copper concentrations (Chapter 5). Therefore, it was not surprising to observe accelerated free
chlorine degradation in the presence of orthophosphate relative to that for corroded copper
coupons alone (single factor ANOVA test, p =0.03).
Corroded lead coupons had comparable overall free chlorine decay constants with the
bulk water control, indicating that bulk water reactions were the primary pathways for chlorine
degradation. There was no significant difference in chlorine decay constants for corroded lead
coupons in the absence or presence of orthophosphate.
Similar to the results observed with Mannheim Water, for corroded metal coupons in
Britannia Water, free chlorine was more reactive with iron, followed by copper and then lead.
Wall reactions dominated over bulk water reactions for corroded iron coupons. However, the
wall consumption of free chlorine was also dominant for copper coupons alone, which was
different from the observations with Mannheim Water. The addition of orthophosphate did not
statistically affect free chlorine decay for corroded iron and copper coupons, but significantly
decreased free chlorine degradation for corroded lead coupons (single factor ANOVA test, p =
0.005).
A two-factor ANOVA test at a confidence level of 95% was also performed to evaluate
the impacts of water quality and the interactive effects of corrosion inhibitors with metal surface
in two water matrices on chlorine decay. The results are shown in Table 3-5. For the three tested
metal materials and bulk water, regardless if orthophosphate was present, free chlorine degraded
more quickly in Mannheim Water than in Britannia Water with all of the p-values below 5%.
This suggests that something specific to the water types (e.g., the relatively higher TOC level in
54
Mannheim Water than in Britannia Water) was the primary reason for the rapid free chlorine
degradation in Mannheim Water.
Table 3-5 Comparison of HOCl overall decay constants for corroded coupons between
Mannheim Water and Britannia Water (n=4)
Material Corrosion
inhibitors
Mannheim
Water
Britannia
Water
p-value (=5%)
Water
quality
Corrosion
inhibitors
Iron None 0.1850±0.0178 0.0953±0.0044
0.0009 0.19 Orthophosphate 0.2064±0.0010 0.1075±0.0242
Copper None 0.0346±0.0056 0.0200±0.0068
0.002 0.08 Orthophosphate 0.0513±0.0028 0.0195±0.0026
Lead None 0.0268±0.0046 0.0156±0.0003
0.009 0.26 Orthophosphate 0.0251±0.0056 0.0081±0.0007
Control
(water)
None 0.0254±0.0007 0.0073±0.0006 0.0002 0.90
Orthophosphate 0.0248±0.0014 0.0070±0.0014
Orthophosphate appeared not to affect free chlorine degradation for the tests performed in
these two tested water matrices since all of the p-values were above 5%. However, by further
comparing chlorine decay constants for corroded copper and lead coupons in the absence and
presence of orthophosphate for each single water matrix (Table 3-5), orthophosphate had a
different impact on free chlorine degradation.
For corroded copper coupons, free chlorine in Mannheim Water degraded more rapidly in the
presence of orthophosphate, but there was no statistical difference in free chlorine degradation in
the absence and presence of orthophosphate for Britannia Water. Figure 3-3 shows copper
release kinetics for corroded copper coupons in two water matrices. In the presence of
orthophosphate, the copper release rate in Mannheim Water was generally increased relative to
that in the absence of orthophosphate. However, orthophosphate slightly decreased the copper
release rate in Britannia Water after 30 hours.
For corroded lead coupons, orthophosphate significantly decreased free chlorine
degradation in Britannia Water (single factor ANOVA test, p = 0.005), but exhibited no
statistical impacts on free chlorine degradation in Mannheim Water. As shown in Figure 3-4,
although orthophosphate effectively reduced lead release in both water matrices, the magnitude
of reduction by orthophosphate was more significant in Britannia Water than that in Mannheim
Water. Therefore, the interactions of orthophosphate and corroded metal surface, which are
55
likely affected by water quality, may play an important role in determining free chlorine
degradation rates for corroded copper and lead coupons in different water matrices.
Figure 3-3 Copper release kinetics for corroded coupons in Mannheim Water and Britannia
Water; error bars indicate the measured maximum and minimum values (n=2)
Figure 3-4 Lead release kinetics for corroded coupons in Mannheim Water and Britannia Water;
error bars indicate the measured maximum and minimum values (n=2)
HAA Formation 3.3.2
3.3.2.1 Fresh Metal Coupons
In this study, 9 compounds of haloacetic acids were analyzed. However, due to relatively
low concentrations of bromodichloro-, chlorodibromo-, and tribromo-acetic acids, only the
concentrations of the other 6 haloacetic acids (HAA6) are reported here. A discussion concerning
the HAA6 speciation will be provided in Section 3.3.3. Figure 3-5 demonstrates formation
kinetics of HAA6 for fresh metal coupons in the absence and presence of corrosion inhibitors in
Mannheim Water. It should be noted that for water samples containing iron coupons, despite no
0.0
0.4
0.8
1.2
0 10 20 30 40 50 60
Cu (
mg/L
)
Time (hours)
Mannheim Water
Cu
Cu+orthophosphate
0.0
0.5
1.0
1.5
2.0
0 10 20 30 40 50 60
Cu (
mg/L
)
Time (hours)
Britannia Water
Cu
Cu+orthophosphate
0
1
2
3
4
0 10 20 30 40 50 60
Pb (
mg/L
)
Time (hours)
Mannheim Water
Pb
Pb+orthophosphate
0
1
2
3
4
0 10 20 30 40 50 60
Pb (
mg/L
)
Time (hours)
Britannia Water
Pb
Pb+orthophosphate
56
chlorine residual after 24 hours, HAA concentrations were monitored to investigate the fate of
HAA6 when they were in contact with iron. As shown in Figure 3-5, the rates of HAA6 formation
for all of the samples were relatively fast in the initial 8 hours, and then tended to level off after
24 hours for iron coupons and after 100 hours for lead coupons and bulk water control. However,
Figure 3-5 HAA6 formations with time in the presence and absence of corrosion inhibitors for
fresh metal coupons with Mannheim Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2)
0
40
80
120
160
200
0 30 60 90 120
HA
A 6
(µ
g/L
)
Time (hours)
CuCu+polyphosphateCu+orthophosphateControl
0
40
80
120
160
200
0 30 60 90 120
HA
A 6
(µ
g/L
)
Time (hours)
Pb
Pb+polyphosphate
Pb+orthophosphate
Control
0
40
80
120
160
200
0 30 60 90 120
HA
A 6
(µ
g/L
)
Time (hours)
FeFe+polyphosphateFe+orthophosphateControl
57
no plateaus of HAA formation were observed for copper coupons in the absence and presence of
corrosion inhibitors, and HAA6 continued to form after 100 hours. Compared with bulk water
control, iron coupons produced approximately 50% less HAA6 at similar reaction time. It is
evident since free chlorine was consumed primarily by iron corrosion and thus less free chlorine
was available to react with HAA6 precursors to form HAA6. No significant influence of
corrosion inhibitors on HAA6 formation was observed for fresh iron coupons.
For water containing copper coupons, lower HAA6 production was also observed at each
reaction time compared with bulk water. However, the HAA6 formation rate for copper alone
was obviously higher than for other copper coupons in the presence of different corrosion
inhibitors and even in the bulk water. When the reaction time was 98 hours, 157 µg/L of HAA6
formed for copper alone, significantly higher than that in the bulk water which was only 141
µg/L (student t test, p-value =0.008). Other water containing copper coupons showed faster
HAA6 formation kinetics than for the bulk water after 72 hours as well, even though the absolute
concentrations of HAA6 were still lower than in the bulk water. The fact that copper corrosion
itself consumed some free chlorine implied that less HAA6 should have been formed in the
system. Enhanced formation of HAA compounds in the presence of higher levels of copper ions
indicated that these copper ions could catalyze HAA formation. The mechanism is believed to be
similar to the catalytic effect of copper(II) and copper oxides in THM formation, in which copper
could complex with THM precursor compounds and enhance oxidative decarboxylation and
enolization of the keto-groups (Blatchley et al., 2003; Li et al., 2008). This new observation also
laid the groundwork for further investigation of copper catalysis on HAA formation in Chapter 5.
Catalytic potential of copper corrosion products on HAA formation has been confirmed, and
copper catalysis was dependent on copper concentration, pH and the types of its solid corrosion
products.
In terms of HAA formation for water containing lead coupons in the absence of corrosion
inhibitors, the consumption of free chlorine by lead corrosion resulted in less HAA6 being
produced. Comparable HAA formation was observed for lead coupons treated with 1 mg/L
polyphosphate and 1 mg/L orthophosphate as bulk water control. Further investigation of the
distribution of HAA6 species (Section 3.3.3) shows that the variation of HAA concentrations
among lead coupons treated with different corrosion inhibitors was primarily due to the
difference in MCAA concentrations that were formed. The MCAA formation rate for lead
58
coupons alone was lower than lead coupons treated with polyphosphate and orthophosphate for
72 hours.
Generally, corrosion inhibitors did not affect HAA formation during short reaction
periods (<12 hours) for fresh metal coupons, which has not been reported previously. However,
no conclusion can be made about the long-term effect of corrosion inhibitors on HAA formation
for water containing iron coupons due to fast depletion of free chlorine within 24 hours. Copper
and lead coupons exhibited different HAA formation rates in the presence of different corrosion
inhibitors. The role of these corrosion inhibitors on HAA formation was essentially due to their
effects on the interactions between free chlorine and metals.
In terms of HAA formation in Britannia Water, Figure 3-6 displays HAA6 formation
kinetics for fresh metal coupons in the absence and presence of corrosion inhibitors and in the
bulk water control. Different from the observations in Mannheim Water in which iron coupons
produced much less HAA6 compared with the bulk water control, iron coupons in Britannia
Water had similar HAA formation kinetics to the bulk water control at least for the first 24 hours.
This indicates that the free chlorine concentration, despite its consumption by iron corrosion,
may not be a rate-limiting factor for HAA formation. The difference in some aspect of the water
quality of Britannia Water from Mannheim Water may also be the reason for the observed
different HAA formation kinetics between these two water matrices. Corrosion inhibitors did not
significantly affect HAA formation in this test water as well.
For fresh copper coupons, in the initial 24 hours, HAA formation kinetics were similar to
those observed in the bulk water. After 24 hours and in the absence of corrosion inhibitors,
HAA6 exhibited a faster formation rate compared with other copper coupons in the presence of
corrosion inhibitors and the bulk water control. This indicates that copper catalysis also played
an important role in the enhanced formation of HAA. However, no differences in HAA
formation rates were observed between copper coupons in contact with corrosion inhibitors and
bulk water control in the duration of tests. In Britannia Water, up to 1.1 mg/L Cu(II) was
released from fresh copper coupons after 24 hours, but only 0.7 and 0.4 mg/L Cu(II) was
released from copper coupons when they were in contact with polyphosphate and
orthophosphate, respectively. This suggests that copper catalysis was a function of dissolved
59
copper concentration, and the catalytic impacts of copper at a concentration below 0.7 mg/L may
not be enough to enhance HAA formation.
Figure 3-6 HAA6 formations with time in the presence and absence of corrosion inhibitors for
fresh metal coupons with Britannia Water in one set of experiments, error bars indicate the
measured maximum and minimum values (n=2)
For the lead coupons, comparable HAA formation rates were observed in the absence and
presence of corrosion inhibitors relative to the bulk water control. Further investigation of the
0
30
60
90
120
150
0 10 20 30
HA
A6 (
µg
/L)
Time (hours)
Fe
Fe+polyphosphate
Fe+orthophosphate
Control
0
30
60
90
120
150
0 20 40 60 80
HA
A6 (
µg
/L)
Time (hours)
Cu
Cu+polyphosphate
Cu+orthophosphate
Control
0
30
60
90
120
150
0 20 40 60 80
HA
A6 (
µg
/L)
Time (hours)
PbPb+polyphosphatePb+orthophosphateControl
60
distribution of HAA6 species showed that MCAA formation rates for all of the lead coupons
were similar in Britannia Water. Along with the observation that higher MCAA formation rates
corresponded to higher HAA formation rates in Mannheim Water, a general conclusion is made
that different HAA6 formation kinetics for lead coupons in the absence and presence of corrosion
inhibitors was likely attributed to their different MCAA formation kinetics.
3.3.2.2 Corroded Metal Coupons
Figure 3-7 displays HAA6 formation at 48 hours for corroded metal coupons in the
absence and presence of orthophosphate compared with the bulk water control. In Mannheim
Water, at 48 hours, corroded iron coupons had less HAA6 formation (40.7~48.7 µg/L) compared
with bulk water (62.3 µg/L). Due to fast consumption of free chlorine by iron corrosion, it is
reasonable to observe reduced formation of HAA relative to bulk water control. For corroded
copper coupons, there was no significant difference in HAA formation compared with that in
bulk water control. Copper has been shown to catalyze HAA formation for fresh copper coupons
(Section 3.3.2.1), and its catalysis increases nonlinearly with increasing dissolved copper
concentrations, and to become more prominent at longer reaction times (Chapter 5). Therefore,
due to the relatively low concentration of the released copper (< 0.5 mg/L) in the solution,
copper catalysis during HAA formation may not be expected to be significant within 48 hours,
Figure 3-7 HAA6 formation at 48 hours in the presence and absence of corrosion inhibitors for
corroded metal coupons with Mannheim Water and Britannia Water in one set of experiments,
error bars indicate the measured maximum and minimum values (n=2)
0
30
60
90
120
HA
A6 (
µg/L
)
Mannheim Water Britannia Water
61
which is consistent with the observed results. Also, since free chlorine degraded at a similar rate
in the bulk water control and the water in contact with lead coupons (Section 3.3.1.2), it was not
surprising to observe similar HAA formation kinetics for corroded lead coupons in Mannheim
Water.
In Britannia Water, HAA formation at 48 hours for all of metal coupons and bulk water
control were statistically the same, regardless of the presence of orthophosphate. As discussed in
Section 3.3.1.2, free chlorine degraded faster in the presence of metal coupons, especially when
in contact with corroded iron coupons. Comparable HAA yields between corroded metal
coupons and the bulk water, again, suggests that free chlorine concentration may not be a rate-
limiting factor for HAA formation in Britannia Water.
3.3.2.3 HAA Formation and Chlorine Demand
Figure 3-8 shows the correlation between consumed free chlorine and HAA6 formation
for bulk water control. Strong correlation between disinfectant demand and HAA6 formation (R2
=0.98) confirms that DBP formation is proportional to disinfectant consumption in the bulk
solution as expected (Clark and Sivaganesan, 1998; Gang et al., 2002).
Figure 3-8 HAA formation and free chlorine demand for bulk water in the absence of corrosion
inhibitors
Although free chlorine was consumed partially as a result of metal corrosion, a similar
linear relationship between the overall free chlorine demand and HAA formation was observed
for all of the iron and lead coupons regardless of metal age, water quality, and the presence of
corrosion inhibitors (plots are not provided). This linear relationship was also observed for
R² = 0.98
0
40
80
120
160
200
0 2 4 6
HA
A6 (
µg
/L)
Chlorine demand (mg/L)
62
corroded copper coupons in the absence and presence of orthophosphate in both water matrices.
It is inferred that metal corrosion products consumed free chlorine linearly as well.
However, an exponential relationship between chlorine demand and HAA6 formation was
observed for fresh copper coupons in the absence and presence of different corrosion inhibitors.
Figure 3-9 shows this nonlinear relationship for copper coupons alone in Mannheim Water. In
the first several hours, HAA formation increased linearly with chlorine demand. It can be
hypothesized that a small amount of copper ions initially released into the water did not affect
HAA formation significantly. When more copper ions were released with time, their catalytic
effects on HAA formation became prominent. More HAA formation was observed after 98 hours
in the presence of copper ions (157 µg/L, Point A) compared with the extrapolated value at 98
hours on the assumed linear line established between initial HAA formation and chlorine
demand (approximately 90 µg/L, Point B). Further experiments have been conducted and the
catalysis of copper of HAA formation has been confirmed (Chapter 5).
Figure 3-9 HAA formation and free chlorine demand for copper in the absence of corrosion
inhibitors
HAA Speciation 3.3.3
3.3.3.1 Fresh Coupons
Figure 3-10 displays the speciation of HAA6 in Mannheim Water when in contact with
fresh metal coupons and in bulk water control as 1 mg/L orthophosphate was applied as a
R² = 1.00
0
50
100
150
200
0 3 6 9 12 15
HA
A6 (
µg
/L)
Chlorine demand (mg/L)
Point A
Point B
Assumed linear HAA formation
63
corrosion inhibitor. Similar results were obtained for Britannia Water, shown in Appendices
Figure 10-15. Since DBAA and MBAA concentrations were consistently low, only DCAA,
TCAA, MCAA and BCAA are shown. Generally, DCAA formation was dominating over the
production of TCAA for all metal coupons, which is consistent with other research (Rossman et
al., 2001). As a result of the reactions between free chlorine and metal coupons, less availability
of free chlorine in the solution limited the formation of TCAA (Rossman et al., 2001).
Figure 3-10 HAA speciation with time in the presence of 1 mg/L orthophosphate for fresh metal
coupons with Mannheim Water in one set of experiments, error bars indicate the measured
maximum and minimum values (n=2)
For iron coupons, the concentrations of DCAA, MCAA and BCAA leveled off after 24
hours. It is primarily because free chlorine was depleted within 24 hours, making one reactant for
HAA formation unavailable. On the other hand, TCAA concentration exhibited the decrease
trend with time after 24 hours. As reported by Hozalski et al. (2001), HAA compounds may
experience reductive transformations by Fe0 via sequential hydrogenolysis (replacement of a
halogen by hydrogen), but no degradation of TCAA was observed over 150 hours in the presence
0
15
30
45
60
75
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
Fe
DCAA TCAA
MCAA BCAA
0
15
30
45
60
75
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
Cu
0
15
30
45
60
75
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
Pb
0
15
30
45
60
75
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
Control
b
64
of goethite, magnetite and aqueous Fe(II) (Chun et al., 2005). XRD analysis was performed on
the surface of iron coupons for short term corrosion, and Fe0 has been identified (Appendices
Figure 10-16). It indicates that 24 hours was not long enough for iron coupons to be fully
oxidized by free chlorine, and the decrease in TCAA was primarily attributed to the reduction by
Fe0 present on the surface of corroded iron coupons. To fully understand the fate of HAA in the
water in contact with iron, experiments have been conducted to investigate the reduction
potential of iron on HAA under controlled experimental conditions, and results are provided in
Appendices 10.9.
For copper coupons, DCAA formation contributed to more than 60% of HAA6 after 24
hours, whereas TCAA, BCAA and MCAA concentrations were generally below 15 µg/L. For
lead coupons, DCAA formation kinetics was slightly faster than TCAA, and up to 22 µg/L of
MCAA formed after 24 hours. The relative contribution of DCAA, TCAA and MCAA to HAA6
was 40%, 30% and 20%, respectively. Similarly, less than 15 µg/L BCAA was formed for lead
coupons during the test period. Bulk water control had comparable DCAA and TCAA formation
rates. MCAA concentrations in bulk water control were consistently lower than DCAA and
TCAA, but higher than BCAA concentrations which were always below 15 µg/L.
3.3.3.2 Corroded Coupons
The observed HAA speciation after 48 hours of contact with chlorine for corroded
coupons in the presence of orthophosphate with Mannheim Water is shown in Figure 3-11.
MBAA, BCAA and DBAA were all <5 µg/L and so they are not shown. MCAA dominated over
the production of DCAA and TCAA for all the investigated metal coupons and the bulk water
control. The dominance of MCAA over DCAA for corroded metal coupons was different from
the observations for fresh metal coupons in the same water matrix. In the tests employing
corroded metal coupons, 5.5 mg/L free chlorine was applied, but a high concentration of free
chlorine (12.3 mg/L) was dosed in the tests employing fresh metal coupons. Therefore, the
difference in HAA speciation between these two tests was likely because the low concentration
of free chlorine applied for corroded metal coupons limited the formation of highly chlorinated
species, such as DCAA and TCAA (Rossman et al., 2001).
65
Figure 3-11 HAA speciation with time in the presence of 1 mg/L orthophosphate for corroded
metal coupons with Mannheim Water in one set of experiments, error bars indicate the measured
maximum and minimum values (n=2)
Figure 3-12 demonstrates HAA speciation at 48 hours for corroded coupons in the
presence of orthophosphate with Britannia Water. DCAA formation was observed to dominate
over TCAA and MCAA formation. A similar trend was also observed for fresh metal coupons in
Britannia Water. The dominance of DCAA formation over other species, regardless of the
concentrations of free chlorine, indicates that free chlorine concentration may not affect relative
distribution of HAA species in Britannia Water.
Figure 3-12 HAA speciation with time in the presence of 1 mg/L orthophosphate for corroded
metal coupons with Britannia Water in one set of experiments, error bars indicate the measured
maximum and minimum values (n=2)
0
10
20
30
40
50
Fe+ortho Cu+ortho Pb+ortho Control
HA
A (
µg
/L)
MCAA DCAA TCAA
0
10
20
30
40
Fe+ortho Cu+ortho Pb+ortho Control
HA
A (
µg
/L)
MCAA DCAA TCAA
66
3.4 Summary
This study compared the kinetics of free chlorine degradation and HAA formation as well
as HAA speciation in two water matrices in contact with three metal materials (either fresh or
aged) commonly found in distribution system water mains (e.g. Fe) and premise plumbing (e.g.
Cu and Pb). MS-FP and MS-SDS tests were conducted on a “survey” basis mainly to help
identify topics/theories for further research (later chapters), and thus no detailed mechanistic
investigation was explored at this point. Most of the results were consistent with those of other
researchers, providing confidence in the testing procedures. New or unexpected results were
either further explored in subsequent tests or simply identified for follow-up by other
researchers.
The following conclusions can be drawn based on the experimental results:
a. In agreement with previous studies, free chlorine degradation generally followed pseudo-
first-order kinetics, regardless of metal age and water quality. The experimental design
uniquely enabled the direct comparison of free chlorine degradation kinetics among three
metal materials. The reactivity of these pipe materials, for both fresh and aged ones, as
investigated in this study was: ductile iron >copper >lead.
b. New observations concerning the impacts of two different water quality on free chlorine
degradation showed that these impacts were dependent on metal type and metal age. The
low pH and alkalinity in Britannia Water led to increased free chlorine decay by
increasing iron corrosion, but only for fresh iron coupons. For all of the corroded metal
coupons, free chlorine degraded consistently more quickly in Mannheim Water than in
Britannia Water, most likely due to the higher TOC level in Mannheim Water.
c. The impacts of phosphate-based corrosion inhibitors on chlorine decay and HAA
formation were newly identified. The addition of these corrosion inhibitors generally
increased HOCl degradation for fresh iron coupons, and they decreased HOCl decay only
for fresh copper coupons. These corrosion inhibitors did not impact HOCl decay for both
fresh and pre-corroded lead coupons. Generally, phosphate-based corrosion inhibitors did
not impact HAA formation, regardless of metal type, metal age and water quality.
d. For fresh copper coupons in both investigated water matrices, HAA formation was
enhanced in the presence of high levels of copper ions, indicating possible catalytic
67
potential of copper on HAA formation. This new observation laid the groundwork for
further investigation of copper catalysis on HAA formation in Chapter 5.
e. Consistent with previous studies, DCAA was the dominant species when water was in
contact with fresh metal coupons. For corroded metal coupons, either MCAA or DCAA
formation was favored.
In addition, although the results from this survey study suggest that metal materials and
their interactions with corrosion inhibitors affected HOCl decay and HAA formation, these
impacts may be site-specific due to their dependence on metal age and water quality. This study
also reveals that the addition of phosphate-based corrosion inhibitors can decrease chlorine
demand, particularly for copper, and thereby allow for the maintenance of high disinfectant
residuals in distribution mains and premise plumbing. Although a higher level of residuals is
desirable to control microbial growth and pathogen amplification in large building such as
hospitals, it may also potentially increase halogenated DBP formation. Therefore, water utilities
may need to decrease their initial chlorine dosage in distribution systems when considering
applying these corrosion inhibitors in their systems.
One should also keep in mind that these tests were conducted for relatively new pipe
materials in bench-scale simulated distribution systems under the static condition. Although pre-
corroded metal coupons were employed in this study, they were aged only for up to one month.
Therefore, all of these results are mainly indicative of what may take place in dead-ends and
during stagnation within plumbing and distribution systems when new pipes are installed and/or
after these new pipes experience only short-term exposure to secondary disinfectants. Any
effects from hydrodynamics, pipe service age and microorganisms should also be taken into
account, and some of these effects were addressed in the pilot-scale experiments that were
performed for Chapter 7. Furthermore, since HAA concentrations may be reduced to meet new
regulations by switching from free chlorine to combined chlorine, but this can increase NDMA
formation, the impacts of copper catalysis on NDMA formation were also examined in
experiments that are summarized in Chapter 6.
68
3.5 References
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Pipe service age effect. Water Research, 41(2), 387-396.
APHA, AWWA, WEF (2005) Standard Methods for the Examination of Water & Wastewater,
21th Edition, Washington D C, USA.
Baribeau, H., Boulos, L., Hileselassie, H., Crozes, G., Singer, P. C., Nichols, C., Schleisinger, S.
A., Gullick, R. W., Williams, S. L., Williams, R. L., Fountleroy, L., Andrews, S. A., and
Moffat, E. (2006) Formation and decay of disinfection byproducts in the distribution
system. Water Research Foundation and US EPA, Project # 2770, Denver, USA.
Blatchley, E. R., Margetas, D., and Duggirala, R. (2003) Copper catalysis in chloroform
formation during water chlorination. Water Research, 37(18), 4385-4394.
Brereton, J. A., and Mavinic, D. S. (2002) Field and material-specific simulated distribution
system testing as aids to understanding trihalomethane formation in distribution systems.
Canadian Journal of Civil Engineering, 29(1), 17-26.
Chen, W.J. and Weisel, C.P. (1998) Halogenated DBP concentrations in a distribution system.
Journal American Water Works Association, 90(4), 151-163.
Chun, C. L., Hozalski, R. M., and Arnold, T. A. (2005( Degradation of drinking water
disinfection byproducts by synthetic goethite and magnetite. Environmental Science &
Technology, 39(21), 8525-8532.
Dartmann, J., Alex, T., Dorsch, T., Schevalje, E., and Johannsen, K. (2004) Influence of
decarbonisation and phosphate dosage on copper corrosion in drinking water systems.
Acta Hydrochimica Et Hydrobiologica, 32(1), 25-32.
Demora, S. J., and Harrison, R. M. (1984) Lead in tap water - contamination and chemistry.
Chemistry in Britain 20(10), 900-904.
Facey, R. M., and Smith, D. W. (1995) Soft, low-temperature water-distribution corrosion:
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69
Fu, J., Qu, J. H., Liu, R. P., Qiang, Z. M., Liu, H. J., and Zhao, X. (2009b) Cu(II)-catalyzed
THM formation during water chlorination and monochloramination: A comparison study.
Journal of Hazardous Materials, 170(1), 58-65.
Gray, E. T., Taylor, R. W., and Margerum, D. W. (1977) Kinetics and Mechanisms of Copper-
Catalyzed Decomposition of Hypochlorite and Hypobromite - Properties of a Dimeric
Copper(III) Hydroxide Intermediate. Inorganic Chemistry, 16(12), 3047-3055.
Hallam, N. B., West, J. R., Forster, C. F., Powell, J. C., and Spencer, I. (2002) The decay of
chlorine associated with the pipe wall in water distribution systems. Water Research,
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Hallam, N. B., Hua, F., West, J. R., Forster, C. F., and Simms, J. (2003) Bulk decay of chlorine
in water distribution systems. Journal of Water Resources Planning and Management-
Asce, 129(1), 78-81.
Haas, C. N., Gupta, M., Chitluru, R., and Burlingame, G. (2002 Chlorine demand in disinfecting
water mains. Journal American Water Works Association, 94(1), 97-102.
Health Canada (2008) Guidelines for Canadian Drinking Water Quality: Guideline Technical
Document — Haloacetic Acids., Water, Air and Climate Change Bureau, Healthy
Environments and Consumer Safety Branch, Health Canada, Ottawa, Ontario.
Health Canada (2009) Guidance on Controlling Corrosion in Drinking Water Distribution
Systems. Water, Air and Climate Change Bureau, Healthy Environments and Consumer
Safety Branch, Health Canada, Ottawa, Ontario (Catalogue No. H128-1/09-595E).
Hozalski, R. M., Zhang, L., and Arnold, W. A. (2001) Reduction of haloacetic acids by Fe0:
Implications for treatment and fate. Environmental Science & Technology, 35(11), 2258-
2263.
Kanokkantapong, V., Marhaba, T. F., Panyapinyophol, B., and Pavasant, P. (2006) FTIR
evaluation of functional groups involved in the formation of haloacetic acids during the
chlorination of raw water. Journal of Hazardous Materials, 136(2), 188-196.
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Leroy, P. (1993) Lead in drinking water: Origins, solubility, treatment. Aqua (Oxford), 42(4),
233-238.
Li, B., Liu, R. P., Liu, H. J., Gu, J. N., and Qu, J. H. (2008) The formation and distribution of
haloacetic acids 9 in copper pipe during chlorination. Journal of Hazardous Materials,
152(1), 250-258.
Li, S. Y., Lixiao, N. X., Sun, C., and Wang, L. S. (2004) Influence of organic matter on
orthophosphate corrosion inhibition for copper pipe in soft water. Corrosion Science,
46(1), 137-145.
Lu, W., Kiene, L., and Levi, Y. (1999) Chlorine demand of biofilms in water distribution
systems. Water Research, 33(3), 827-835.
MacQuarrie, D. M., Mavinic, D. S., and Neden, D. G. (1997) Greater Vancouver Water District
drinking water corrosion inhibitor testing. Canadian Journal of Civil Engineering, 24(1),
34-52.
Maddison, L.A. and Gagnon, G.A. (1999) Evaluating corrosion control strategies for a pilot scale
distribution system. In Proceedings of the 1999 American Water Works Association
Annual Conference, Tampa Bay, Fla. American Water Works Association, Denver, Colo.
McNeill, L. S., and Edwards, M. (2000) Phosphate inhibitors and red water in stagnant iron
pipes. Journal of Environmental Engineering-Asce, 126(12), 1096-1102.
Montgomery, D. C. (2000) Design and Analysis of Experiments, 5th edition, John Wiley &
Sons, New York.
Powell, J. C., Hallam, N. B., West, J. R., Forster, C. F., and Simms, J. (2000) Factors which
control bulk chlorine decay rates. Water Research, 34(1), 117-126.
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71
Rossman, L. A., Brown, R. A., Singer, P. C., and Nuckols, J. R. (2001) DBP formation kinetics
in a simulated distribution system. Water Research, 35(14), 3483-3489.
Schock M, Wagner I. (1985) The corrosion and solubility of lead in drinking water. In: Internal
Corrosion of Water Distribution Systems. Denver: AWWA Research Foundation,
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Scholze R.J.; Pontow K.A.; Kanchlbhatla G. and Ray B.T. (1994) Using the CERL pipe loops
system (PLS) to evaluate corrosion inhibitors that can reduce lead in drinking water
(FEAP-TR-EP-94/04). U.S. Army Construction Engineering Research Laboratories,
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Singley, J. E. (1994) Electrochemical nature of lead contamination. Journal American Water
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72
A Comparison of Iron, Copper and Lead Corrosion in 4
Simulated Distribution Systems
Experiments that formed the basis for Chapters 3 and 4 compared the kinetics of metal
corrosion, chlorine degradation and HAA formation for iron, copper and lead treated with
different corrosion inhibitors on a “survey” basis. All of the metal release kinetics data were
collected at the same time as for Chapter 3, but are discussed as a separate chapter herein
because of the quantity of the data. Most of the observations regarding the impacts of corrosion
inhibitors, water quality, and disinfectant type follow the trends that have been reported by other
researchers, and thus no mechanistic investigations were involved in this chapter. Interpretations
concerning the impacts of these parameters on metal corrosion essentially provided valid
explanations for the findings in Chapter 3.
As such, the results in this chapter that deal with metal release kinetics for fresh metal
coupons in one type of water (Mannheim Water, Section 4.3.1) have been published as part of:
Zhang, H., Andrews, S. A. (2012) Effects of Phosphate-based Corrosion Inhibitors on the
Kinetics of Chlorine Degradation and HAA Formation in Contact with Three Metal
Materials. Canadian Journal of Civil Engineering, 39, 44-54.
Data from similar experiments using another water matrix (Britannia Water) and experiments
using pre-corroded coupons for both tested water matrices are included in this chapter for
completeness.
Generally, results from this chapter focus on the research gap "Do metal age and water
quality affect the efficacy of corrosion inhibitors?"
73
Abstract
Pipe materials and their corrosion products may affect the formation and the degradation
of disinfection byproducts in distribution systems. The investigation of factors influencing metal
corrosion under controlled experimental conditions is important for evaluating the impacts of
metal type, corrosion inhibitors and water quality on disinfectant degradation and DBP
formation. This study examined the effects of disinfectant type, metal age, phosphate-based
inhibitors, and water quality on metal corrosion at bench scale. Orthophosphate had no beneficial
effects to control iron corrosion, regardless of water quality, disinfectant type or the age of the
metal surface. However, orthophosphate significantly decreased released copper concentrations,
in particular for fresh coupons during short term exposures, and it also significantly reduced lead
release, irrespective of the age of the metal surface or water quality. In two investigated water
matrices, the water with lower pH and alkalinity generally exhibited higher corrosion potential to
metal coupons. Only for pre-corroded copper coupons, HOCl was consistently more aggressive
than NH2Cl. The results of XPS surface analysis suggest that the effectiveness of orthophosphate
on copper and lead corrosion control was due to facilitated precipitation of calcium. For iron
coupons, CaCO3 precipitation decreased the amount of iron oxides formed on the surface.
Keywords: Ductile iron; Copper; Lead; Corrosion inhibitor; Secondary disinfectant; XPS
74
4.1 Introduction
Corrosion is defined as “the deterioration of a material that results from a reaction with
its environment” (NACE International, 2009). In drinking water distribution systems, iron pipes
are widely used, while in household plumbing systems copper pipes are commonly found.
Although the use of lead in drinking water applications has been eliminated in Canada since
1992, lead may be leached into drinking water from old lead service lines, soldered joints and
brass plumbing fittings (Health Canada, 2009). The established drinking water guideline for lead
based on its health effects in children is 0.010 mg/L. Guidelines for iron and copper based on
aesthetic considerations are ≤0.3 mg/L for iron and ≤1.0 mg/L for copper (Health Canada, 2009).
Metal corrosion in distribution systems is affected by many physical and chemical
factors. Among water quality parameters, pH and alkalinity are the most significant factors
influencing metal levels in distribution systems. It has been consistently reported that the
solubility of lead, copper and iron decreases with elevated pH (McNeill and Edwards, 2001;
Sarin et al., 2003; Schock, 1989; Xiao et al., 2007). Normally, lower iron corrosion rates and
iron concentrations in distribution systems have been associated with higher alkalinity due to the
formation of less soluble siderite (FeCO3) (Sarin et al., 2003). In contrast, copper corrosion
release was increased at higher alkalinity due to the formation of soluble cupric bicarbonate and
carbonate complexes (Edwards et al., 2002; Edwards et al., 1996; Schock and Lytle, 1995). The
degree to which alkalinity affects lead solubility depends on the water pH and the form of lead
carbonate present on the pipe surface (Schock, 1980; Schock, 1990).
Free chlorine (HOCl) has been observed to increase metal corrosion rates due to its high
oxidation potential relative to monochloramine (NH2Cl) (Boulay and Edwards, 2001; Cantor et
al., 2003; LeChevallier et al., 1990a; Rahman et al., 2007; Vasquez et al., 2006). However,
ammonia (a product from chloramines hydrolysis) may increase the corrosivity of water on
copper tubing by forming stable copper-ammonia complexes (Boyd et al., 2008; Schock and
Lytle, 1995).
Addition of phosphate-based corrosion inhibitors is an effective strategy to control
corrosion. The basis for orthophosphate addition as a method of corrosion control relies on the
fact that impervious solid films will be developed as a barrier between the corroding metal and
75
the water (Demora and Harrison, 1984; Singley, 1994), whereas polyphosphates are strong
chelating agents which have been successfully used for reducing scaling and sequestrating Fe2+
ions to treat “red water” (Facey and Smith, 1995; Maddison et al., 2001; Williams, 1990).
Although orthophosphate may reduce copper release in the short term by forming a cupric
phosphate scale, in the long term it may interfere with the natural evolution of scales to form the
least soluble species, malachite, leading to increased copper concentrations in the water (Cantor
et al., 2000). Some studies have also reported detrimental effects of polyphosphate on lead
corrosion control due to the formation of lead polyphosphate complexes (Leroy, 1993;
MacQuarrie et al., 1997).
Although considerable studies have been focused on understanding the effects of
corrosion inhibitors and disinfectant type on metal corrosion, most of these studies were
performed in pilot-scale and full-scale distribution systems with aged metal pipes. In these
studies, the impacts of corrosion inhibitors, disinfectant type and water quality on metal
corrosion may have been somewhat confounded by hydrodynamics and microorganisms.
Therefore, the objective of the current study was to assess the effects of phosphate-based
inhibitors, disinfectant type and water quality on metal corrosion for different metal ages under
controlled experimental conditions at bench scale. HOCl and NH2Cl were applied as secondary
disinfectants, and two water matrices with distinctly different water quality were examined in
regards to their corrosion potential. Metal coupons of ductile iron, copper and lead were
employed. These pipe materials are widely present in either water mains (such as ductile iron) or
household plumbing systems (such as copper), and lead may be leached into drinking water from
old lead service lines, soldered joints and brass plumbing fittings (Health Canada, 2009). Results
of this study will help utilities and households identify possible reasons for increased corrosion
rates to aid in taking effective measures to reduce corrosion in their systems.
4.2 Materials and Methods
Reagents and Materials 4.2.1
All chemicals used in this study were ACS grade or higher. The chlorine dosing solution
(approximately 3500 mg/L as Cl2) was prepared by diluting a concentrated solution of sodium
hypochlorite (NaOCl, 6%, VWR) in Milli-Q water. The NH2Cl dosing solution was preformed
76
by adding the free chlorine dosing solution to a 50 mmol/L ammonium chloride solution in a
well-stirred 250 mL amber bottle at a Cl2/N molar ratio of 0.8:1. The target disinfectant
concentrations applied in the tests were achieved by spiking the disinfectant dosing solution into
1 L test water. An initial disinfectant concentration of 12.3 mg/L as Cl2 was applied in
experiments with fresh coupons, and an initial disinfectant concentration of 5.5 mg/L as Cl2 was
applied for experiments employing corroded metal coupons. Orthophosphate (Na3PO4) and
polyphosphate ([Na(PO2)]6) were selected as corrosion inhibitors to investigate their impacts on
metal corrosion. They were prepared at a concentration of 500 mg/L as P and were kept in the
dark at 4ºC. The targeted dosage of both corrosion inhibitors was 1 mg/L as P.
Test coupons of ductile iron, copper and lead were purchased from Metal Samples Co.,
Alabama, US. The size of these coupons was 1 2”3”1 1 ”. For experiments with fresh
coupons, before each batch of experiments, any corrosion products were removed from the
coupons with 60-grit sandpaper followed by 120-grit sandpaper, and then rinsed with deionized
water and acetone followed by Milli-Q water (Scholze et al., 1994). Polished coupons were
submerged into the test water to simulate contact with new pipes without any impacts from
service age. In addition, to investigate the aged pipe environment, some coupons of each pipe
material were pre-conditioned in tap water in the absence and presence of orthophosphate. The
initial free chlorine concentration for conditioning purposes was 10 mg/L. After each 24 hour
reaction period for the iron coupons and each 48 hour reaction period for the copper and lead
coupons, the free chlorine residual was measured and then the water was replaced with fresh tap
water dosed with 10 mg/L free chlorine. The conditioning process proceeded until the 24 hour
chlorine residual for each combination of metal type and corrosion inhibitors was approximately
constant (< 4.0 mg/L chlorine decay for the iron coupons alone, 6.8 mg/L for the iron coupons
treated with orthophosphate, 2.7 mg/L for all of the copper coupons, and 1.9 mg/L for all of the
lead coupons). The measured 24 hour chlorine residuals for iron, copper and lead are shown in
Appendices Figure 10-14. Orthophosphate was selected since it can effectively control metal
release for fresh metal coupons in a short time, especially for copper and lead.
Unchlorinated post-filtration water was collected from two water treatment plants in
Ontario (Mannheim Water and Britannia Water). Water quality parameters are listed in Table
4-1.
77
Table 4-1 Summary of water quality parameters for two post-filtration water sources
Parameters Mannheim Water Britannia Water
pH 7.5 ± 0.2 6.6±0.5
Alkalinity (mg/L) 187±33 12±2
UV254 (cm-1
) 0.058±0.008 0.065±0.005
TOC (mg/L) 4.5±0.3 3.3±0.2
SUVA (L/mg·cm-1) 0.013±0.002 0.019±0.0002
Bromide (μg L) 65.0±15.5 15.8±0.2
Chloride (mg/L) 84.5±2.5 3.1±0.6
Sulfate (mg/L) 35.0±2.0 23.8±0.8
Cl-:SO4
2- ratio 2.4±0.1 0.13±0.04
Experimental Procedures 4.2.2
“Material-specific” simulated distribution system (MS-SDS) and material-specific
formation potential (MS-FP) tests were applied in this study. Details of MS-SDS procedures
have been described by Brereton and Mavinic (2002). In experiments with corroded metal
coupons, MS-SDS tests were performed. In experiments with fresh metal coupons, MS-FP tests
were carried out similarly to MS-SDS tests, except that MS-FP tests applied a higher
concentration of initial disinfectant (12.3 mg/L) than would be encountered in distribution
systems to meet the high chlorine demand of the fresh metal coupons and ensure detectable
disinfectant residuals (0.2 mg/L) after 24 hours.
All of the MS-SDS and MS-FP experiments were performed in 1 L amber bottles with
PTFE lined caps and at room temperature (21±2 ºC). The reaction bottles were made chlorine
demand free before use. Metal coupons of single metal species were suspended in the bottles
using nylon threads. The testing of galvanic effects resulting from combinations of different
metal coupons was beyond the scope of this investigation. To maximize the contact of water with
coupon surface, a Big Bill Orbital shaker was used (Barnstead International) to maintain gentle
mixing at a speed of 25 rpm. After each designated reaction time, 50 mL of water were
withdrawn from the test bottles for the analysis of disinfectant residual and metal concentrations.
Seven to ten time points were employed for each of the kinetic studies. As will be discussed in
Section 4.3.1, metal concentrations reached equilibrium after 10 hours, and the reduction of
water volume due to sequential sampling did not significantly affect equilibrated metal
concentrations. In addition, since all of the tests were conducted in closed systems, the impact of
78
oxygen on metal corrosion was considered to be negligible and thus its concentration was not
measured in this study.
Two sets of kinetic experiments were performed to test the reproducibility of the results,
and duplicate tests were conducted in each set of experiments. All of the tests also included
control samples, which were prepared the same as test samples but without metal coupons.
Analytical methods employed to examine analytes of interest are summarized in Table 4-2.
Table 4-2 Summary of analytical methods
Analyte Unit Instrument /procedure Reference method
TOC mg/L O1-Analytical TOC analyzer SM1 5310 C
pH VWR Scientific Model 8015 pH
meter
UVA254 cm-1
Hewlett Packard 8452A Diode
Array UV spectrophotometer SM 5910B
Alkalinity mg/L Titration SM 2320B
Chlorine mg/L Hach DR2700 Spectrophotometer SM 4500-Cl G
Monochloramine mg/L Hach DR2700 Spectrophotometer Hach method 101712
Anions mg/L Dionex DX-300 Series Ion
Chromatography System SM 4110 B
Metal (Fe, Cu and Pb) mg/L Varian SpectrAA.20 SM 3111 B
Note: 1. SM represents Standard Methods for the Examination of Water and Wastewater APHA et al., 2005. 2. Hach, 2007;
Coupon Surface Analysis 4.2.3
Metal surface analysis, including X-ray diffraction (XRD) and X-ray photoelectron
spectroscopy (XPS), were performed on metal surfaces to identify corrosion products which may
govern the solubility of metal in the solution. For fresh metal coupons, XRD analysis was
performed at the University of Toronto (Department of Chemistry) to identify solid-phase
corrosion products. XRD is a versatile, non-destructive analytical technique that reveals the
mineralogy of the dominant scale solids. For iron coupons, non-uniform corrosion layers were
formed on the coupon surface as a result of short-term reactions, and thus the XRD analysis was
carried out on powders which were scratched from the coupon surface and put on low-
background silicon sample holders (Tang et al., 2006). For lead and copper coupons, the analysis
was performed by placing the samples directly in the diffractometer. A Siemens D5000
Diffractometer System operating at 50 kV/35 mA was used to collect the diffraction patterns. A
high-power, line focus Cu-K-source was used combined with a solid state Kevex detector for
79
elimination of K-lines. The experimental data were collected on a step scan mode (0.02° /1.5
sec) within the most informative range (2-theta degrees). The obtained data were processed by
various Diffrac Plus software packages including Eva 8.0 and Topas v. 2.1.
Surface analysis of corroded metal coupons was performed with XPS at the University of
Toronto (Department of Chemical Engineering). Before XPS analysis, corroded metal coupons
were stored in an anaerobic glove box to remove adsorbed moisture. A Thermo Scientific K-
Alpha XPS system (East Grinstead, UK) was used to scan all of the metal samples. A
monochromatic Al K-Alpha X-ray source was used. The vacuum pressure was approximately
1x10-7
mbar. The survey spectra were acquired with high pass energy (200 eV) and low point
density (1 eV step size). The regional data, used for quantitative evaluation and fitting, was
acquired with a pass energy of 50 eV and a high point density (0.1 eV). The data acquired from
the instrument was processed using the software Avantage (provided by the manufacturer). Any
charging shift produced by the samples was carefully removed by using a 284.6 eV adventitious
carbon (C1s) standard.
4.3 Results and Discussion
Metal Release and Phosphate-based Corrosion Inhibitors 4.3.1
4.3.1.1 Fresh Coupons in Mannheim Water
Iron, copper and lead release kinetics in the absence and presence of corrosion inhibitors
under chlorination in Mannheim Water is illustrated in Figure 4-1. Iron, copper and lead release
kinetics under chloramination exhibited a similar pattern with and without the addition of
corrosion inhibitors, the data for which are shown in Appendices Figure 10-17. Generally, metal
concentrations increased rapidly with increasing reaction time initially, and then more gradually
increased or stabilized after 10 hours. Lower dissolved metal concentrations were observed for
copper and lead coupons when in the presence of the corrosion inhibitors. However, for iron
coupons, phosphate-based corrosion inhibitors did not reduce the amount of iron release under
both chlorination and chloramination, which was also evidenced by McNeill and Edwards (2000
and 2001).
80
Figure 4-1 Kinetics of metal release from fresh metal coupons in the presence and absence of
corrosion inhibitors with HOCl in Mannheim Water; disinfectant concentrations, 12.3 mg/L;
error bars indicate the measured maximum and minimum values (n=2)
The X-ray diffraction patterns of iron powders obtained from the scratched surface
substances of freshly oxidized iron coupons in the absence and presence of different corrosion
inhibitors demonstrate almost identical diffraction patterns for these iron coupons (Figure 4-2).
This may explain the similar metal release kinetics for fresh iron coupons regardless of the
presence of corrosion inhibitors. The main component on the surface of three oxidized iron
coupons was identified as poorly crystalline goethite, α-FeOOH.
0
2
4
6
8
10
0 5 10 15 20 25 30F
e (
mg/L
) Time (hours)
HOCl
FeFe+polyphosphateFe+orthophosphate
0
1
2
3
4
0 20 40 60 80 100 120
Cu (
mg/L
)
Time (hours)
Cu
Cu+polyphosphate
Cu+orthophosphate
0
1
2
3
4
0 20 40 60 80 100 120
Pb (
mg/L
)
Time (hours)
Pb
Pb+polyphosphate
Pb+orthophosphate
81
Figure 4-2 Comparison of XRD patterns for iron powders scratched from the surface of oxidized
iron coupons in the absence and presence of corrosion inhibitors
For fresh copper coupons, phosphate-based corrosion inhibitors significantly reduced
copper release for this water matrix during the short-term kinetic studies in comparison with
copper coupons alone (Figure 4-1), which is consistent with previous studies. Orthophosphate at
1 mg/L reduced copper concentration more effectively than polyphosphate, with the copper
concentration generally below 0.52 mg/L. Results of the XRD analysis demonstrate that only
Cu2O was present on the surface of tested coupons. Figure 4-3 illustrates the relative distribution
of Cu2O on the surface of tested coupons in the absence and presence of corrosion inhibitors. A
smaller percentage of Cu2O (2.25%) in the presence of orthophosphate suggested that
orthophosphate could block active sites on the metal surface, and consequently protect metallic
copper against oxidation (Dartmann et al., 2004). In the presence of polyphosphate, relatively
higher concentrations of copper (approximately 1.1 mg/L) were detected compared with the
orthophosphate treated coupons, but these copper concentrations were consistently lower than
those for copper coupons alone (up to 2.5 mg/L).
The reduction of lead release in the presence of 1 mg/L orthophosphate was also
observed for both HOCl and NH2Cl in Mannheim Water. In contrast, polyphosphate significantly
increased lead solubility, especially when HOCl was used (Figure 4-1). Surface analysis of lead
coupons by XRD indicated that the surface of these coupons had very similar compositions
(Figure 4-4). Aside from the large reflections from the lead, small peaks of lead oxide, litharge
(PbO), could be identified. While hydrocerussite, Pb3(CO3)2(H2O)2, was also possibly present,
PbO was the dominant species.
α-FeO(OH)
α-FeO(OH)
α-FeO(OH)
Fe+poly
Fe+ortho
Fe
10 20 30 40 50 60 70 80
2-theta
Inte
nsity
— Reference pattern of α-FeO(OH)
82
Figure 4-3 Comparison of XRD patterns for copper coupons in the absence and presence of
corrosion inhibitors.
Figure 4-4 Comparison of XRD patterns for lead coupons in the absence and presence of 1 mg/L
orthophosphate
For all of the investigated metal coupons, no phosphate-containing solids were identified
on the surface of coupons exposed to the phosphate-spiked water. Since these tests were initiated
by submerging fresh metal coupons into the test water immediately after free chlorine addition, it
is possible that these phosphate-containing solids were poorly crystalized within 100 hour
reaction time so as not to be identified by XRD (Hsu, 1982; Moriarty, 1990). Nevertheless,
phosphate corrosion inhibitors may still form an amorphous or semi-amorphous protective film
15 20 30 40 50 60 2-theta
Pb+ortho Pb
Pb
O
Pb
O
Pb Pb Pb
Pb
3(C
O3) 2
(OH
)
Inte
nsity
— Reference pattern of Pb; — Reference pattern of PbO; — Reference pattern of Pb3(CO3)2(OH)2
Pb
O
Sample Cu, %. Cu2O, % Cu-control 100.00 0 Cu+ortho 97.75 2.25 Cu+poly 97.07 2.93 Cu 94.58 5.42
Cu Cu+poly Cu+ortho Cu+poly/ortho
Cu control
Cu2O
Cu Cu
Cu2O
28 30 40 50 60
2-theta (degree)
— Reference pattern of Cu2O
Inte
nsity
83
and provide a barrier between metal surface and chlorinated water, thereby reducing the released
metal concentrations. The addition of phosphate corrosion inhibitors may also facilitate the
precipitation of already released metal ions, as insoluble/low soluble metal salts (such as copper
phosphate and lead phosphate), thereby causing the measured metal concentrations to be lower in
the presence of phosphate corrosion inhibitors than those in the absence of corrosion inhibitors
(McNeill and Edwards, 2004). In the case of polyphosphate-treated lead coupons, the increase in
lead concentrations relative to lead coupons alone (especially under chlorination) can be
explained such that polyphosphate increased the solubility of already released lead ions by
complexation (Edwards and McNeill, 2002; Leroy, 1993; MacQuarrie et al., 1997).
4.3.1.2 Corroded Coupons in Mannheim Water
Metal release kinetics for corroded iron, copper and lead coupons was also investigated.
The released metal concentrations following 24 hour exposure under chlorination and
chloramination in Mannheim Water are illustrated in Figure 4-5. Irrespective of the type of
disinfectant, orthophosphate at 1 mg/L increased iron release by at least 90% compared with that
in the absence of orthophosphate. This is likely because orthophosphate increased the solubility
of iron by forming iron-phosphate complexes, as evidenced by McNeill and Edwards (2000). For
corroded copper coupons under chlorination, orthophosphate increased copper release by
approximately 24%, which is different from the observations in the experiments employing fresh
coupons (Figure 4-1). However, this agrees with previous reports in which orthophosphate
interfered with the natural evolution of scales, reducing the formation of the least soluble species
(malachite) and thus having an adverse impact on cuprosolvency for aged pipes, especially in
chlorinated distribution systems (Cantor et al., 2000; Schock and Lytle, 1995; Schock and
Sandvig, 2009). Under chloramination, however, orthophosphate was still effective to reduce
copper concentration. In terms of corroded lead coupons, orthophosphate effectively minimized
lead corrosion irrespective of disinfectant type. Lead concentrations in the presence of
orthophosphate were consistently below 0.1 mg/L under both conditions.
The impacts of phosphate-based corrosion inhibitors on iron and lead release in Britannia
Water were similar to the trends observed with Mannheim Water that orthophosphate had no
beneficial effects on iron release but effectively suppressed released lead concentrations for both
84
Figure 4-5 Metal concentrations after 24 hours for corroded coupons in the absence/presence of
orthophosphate with HOCl and NH2Cl in Mannheim Water, initial disinfectant concentrations
5.5 mg/L; error bars indicate the measured maximum and minimum values (n=2)
fresh and corroded coupons regardless of disinfectant type (Appendices, Figure 10-18 and Figure
10-19). Interestingly, orthophosphate also effectively suppressed copper release for both fresh
and corroded copper coupons under chlorination and chloramination in Britannia Water, while in
Mannheim Water orthophosphate increased the copper corrosion rate for corroded copper
coupons under chlorination (Figure 4-5). The different impacts of orthophosphate on corroded
copper corrosion in the two water matrices indicate that some aspect of the water quality affects
the effectiveness of orthophosphate to control copper corrosion. The dependency of
orthophosphate on water quality to control copper corrosion has also been reported by Dartmann
et al. (2004).
XPS Results for Corroded Coupons in Mannheim Water 4.3.2
Surface analysis by X-ray photoelectron spectroscopy (XPS) was performed on corroded
metal coupons after kinetics experiments with Mannheim Water to identify the compositions of
the corrosion products on the surfaces of the metal coupons and to evaluate the impacts of
disinfectant type and orthophosphate on metal corrosion. Survey scans provided an elemental
analysis of each sample, and high resolution scans allowed specific products to be identified.
0.0
0.2
0.4
0.6
0.8
Me
tal co
nce
ntr
atio
n (
mg
/L)
HOCl NH2Cl
85
4.3.2.1 Survey Scans
Through the survey scan, the elements of carbon, calcium and oxygen were detected on
the surface of most of the corroded metal coupons. Phosphorus was detected only on the surface
of coupons exposed to the orthophosphate-spiked water. For corroded copper and lead coupons,
the percentage of calcium increased significantly when orthophosphate was applied, indicating
that the application of orthophosphate facilitated the precipitation of calcium. The similar
observations were also reported by Stone (2008). XPS survey scan results for corroded copper,
iron and lead coupons under chloramination are displayed in Figure 4-6. The XPS survey scans
for three types of corroded metal coupons under chlorination are displayed in Appendices Figure
10-20.
As shown in Figure 4-6, calcium on the surface of corroded copper coupons increased
from 0.5% in the absence of orthophosphate to 2% when orthophosphate was applied. For lead,
6% of calcium was detected for lead coupons exposed to the orthophosphate-spiked water, but
no calcium was found on the surface of lead coupons in the absence of orthophosphate. The
binding energy of the calcium peak (347.2~347.4 eV) indicated that calcium was present
primarily in combination with phosphate. The formation of calcium phosphate may function as
an additional protective layer to decrease the exposure of the copper surface to the disinfectant
residual, leading to lower levels of copper released into the test water (as shown in Figure 4-5).
However, for corroded iron coupons, calcium did not vary significantly before and after the
addition of orthophosphate, indicating that there was no enhanced precipitation of calcium by
phosphate on the surface of the corroded iron coupons. In this case, the detected phosphorous
was likely present primarily in the form of iron-phosphate complexes since phosphate anions
have a high affinity for iron (Persson et al., 1996). As a result, these complexes would increase
the solubility of iron in the water and result in elevated iron concentrations when orthophosphate
was applied (Figure 4-5).
86
Figure 4-6 Comparison of elemental distribution for copper, lead and iron coupons in the
absence and presence of orthophosphate with NH2Cl
4.3.2.2 High-Resolution Scans
Identification of corrosion products were performed through deconvolutions of high-
resolution XPS scans for the corroded iron, copper and lead coupons (Figure 4-7). For corroded
44%
2% 5%
42%
7%
Cu+NH2Cl+orthophosphate
C
Ca
Cu
O
P
52%
0.5 6%
41%
Cu+NH2Cl
C
Ca
Cu
O
58% 17%
25%
Pb+NH2Cl
Pb
C
O
7%
11%
58%
8%
6%
10%
Pb+ortho+NH2Cl
Pb
C
O
Al
Ca
P
21%
56%
0.9
22%
Fe+NH2Cl
Fe
O
Ca
C
17%
51%
1.2
24%
7%
Fe+ortho+NH2Cl
Fe
O
Ca
C
P
87
Figure 4-7 Corrosion products of Fe, Cu and Pb as well as their relative distribution in the scales
iron coupons, magnetite (Fe3O4) and hydrated ferric oxide (FeOOH) were detected in the scales
regardless of the presence of orthophosphate and disinfectant type. Binding energies for Fe3O4
and FeOOH were 710~710.2 eV and 711.4~711.7 eV, respectively. Both compounds account for
similar fractions in the corrosion scales for all four of the investigated conditions. Under
chlorination, more iron was oxidized in the presence of orthophosphate than in the absence of
0
2
4
6
8
0
2
4
6
8
10
HOCl HOCl+ortho NH2Cl NH2Cl+ortho
Calc
ium
(%
)
Fe,
%
Fe3O4 FeOOH Calcium
0
2
4
6
8
10
HOCl HOCl+ortho NH2Cl NH2Cl+ortho
Cu,
%
Cu2O CuO Cu(OH)2
0
10
20
30
40
50
HOCl HOCl+ortho NH2Cl NH2Cl+ortho
Pb,
%
Pb PbO2 PbO Pb3(OH)2(CO3)2
88
orthophosphate (4.4% vs. 2.6%), indicating that orthophosphate did not protect the metal surface
against oxidation and thus was not beneficial for iron corrosion control. However, under
chloramination, the addition of orthophosphate decreased the amount of oxidized iron in the
scales. Therefore, although orthophosphate was observed to be detrimental for iron release in the
two water matrices investigated in this study, its effects on solid corrosion scales under HOCl
and NH2Cl were inconclusive.
In addition, it was interesting to observe that iron oxides in the scales of iron coupons
under chlorination were less abundant than those under chloramination, and that the sequence of
the amount of oxidized iron in the corrosion scales for four investigated situations was opposite
to the trend in the amount of calcium in the scales. The binding energy of calcium was
346.6~347eV, indicating that it was present mainly in the form of CaCO3. Therefore, CaCO3
played a significant role in determining the abundance of oxidized iron in the corrosion scales,
which has not been reported previously. It followed that greater amounts of CaCO3 precipitation
lead to smaller amounts of iron oxides formed on the metal surface.
Copper corrosion products on the corroded copper coupon surface included Cu2O
(932.2~932.5 eV), CuO (934.0~934.3 eV), Cu(OH)2 and/or CuCO3 (935.2~935.4 eV). Copper
orthophosphate solids, such as cupric phosphate dihydride [Cu3(PO4)2·2H2O] and cupric
phosphate [Cu3(PO4)2], were also possibly present on the surface of copper coupons exposed to
orthophosphate, but they were not identified conclusively. For all of the investigated conditions,
Cu2O was the most abundant species compared with other copper corrosion products, and thus
was the controlling solid phase for copper solubility. In addition, relatively greater amounts of
oxidized copper were detected for coupons exposed to HOCl than to NH2Cl. For example, in the
absence of orthophosphate, 6.7~9.0 % of copper was present in the scale of copper coupons
under chlorination, but only 3% was present under chloramination. This was primarily because
the high oxidation potential of HOCl accelerated copper corrosion rates (Schock and Lytle,
1995). The addition of orthophosphate for both HOCl and NH2Cl tended to decrease the
oxidation rate of Cu(0) to Cu(I) and Cu(II). All this is in agreement with previous studies in
which orthophosphate was able to form an insoluble film on the metal surface to prevent copper
from oxidation, thereby reducing copper release (Edwards et al., 2002).
89
For corroded lead coupons, PbO (litharge, 139~139.1 eV), Pb3(OH)2(CO3)2
(hydrocerussite, 138.1 eV), PbCO3 (cerussite, 138.4 eV), and PbO2 (plattnerite, 137.1~137.4 eV)
were identified as possible corrosion products. The formation and transformation of these lead
species changed with disinfectant type and the presence/absence of orthophosphate. Generally,
oxidized lead was present in the scales of lead coupons in the presence of orthophosphate but
was less abundant than on coupons in the absence of orthophosphate. Only 3.1% and 3.3% of
lead was in the oxidized form in the presence of orthophosphate under chlorination and
chloramination, respectively, but 25% and 43.7% were oxidized in the absence of
orthophosphate with free chlorine and chloramine, respectively. This was most likely because the
facilitated precipitation of calcium by orthophosphate formed a solid Ca3PO4 film and thus
protected lead against further oxidation, which is also evidenced by lower concentrations of lead
in the presence of orthophosphate in Figure 4-5. In addition, PbO2 was detected as a predominant
oxide for lead coupons without the addition of orthophosphate under chlorination (Figure 4-7).
This is mainly because HOCl is a strong oxidant which can oxidize lead to a high oxidation state
(+4), and orthophosphate inhibits the oxidation of Pb(II) to Pb(IV) by forming an insoluble scale
on the metal surface. It is reasonable to find a small amount of Pb3(OH)2(CO3)2 (hydrocerussite)
in the corrosion scales for lead coupons in contact with HOCl since PbO2 can be formed through
further oxidation of Pb3(OH)2(CO3)2 by HOCl (Kim and Herrera, 2010). For lead coupons
exposed to orthophosphate under chlorination and lead coupons alone under chloramination,
Pb3(OH)2(CO3)2 (hydrocerussite) and/or PbCO3 (cerussite) were dominant species, whereas PbO
was a solid controlling phase for lead coupons in the presence of orthophosphate under
chloramination.
Water Quality and Disinfectant Type 4.3.3
The ability of orthophosphate to mitigate impacts of water quality parameters (pH,
alkalinity and NOM) and disinfectant type (HOCl vs NH2Cl) was tested with fresh and pre-
corroded coupons and two water types. As shown in Table 4-1, Britannia Water had a lower pH
(6.60.5) and a significantly lower level of alkalinity (12 mg/L as CaCO3) than Mannheim
Water, pH and alkalinity of which were 7.5 ± 0.2 and 187 mg/L as CaCO3. Furthermore, HOCl
has a higher oxidation potential than NH2Cl. Therefore, it was hypothesized that Britannia Water
would be more corrosive than Mannheim Water, and more metal ions would be released under
chlorination than under chloramination. To test these hypotheses, the released metal
90
concentrations that were measured after exposure to similar initial concentrations of HOCl and
NH2Cl were compared. The results are displayed in Figure 4-8.
Figure 4-8 Comparison of iron concentrations at 24 hours in the presence of orthophosphate for
fresh and corroded coupons under free chlorine and chloramine, error bars indicate the measured
maximum and minimum values (n=2)
0
2
4
6
8
10
MannheimWater:fresh
MannheimWater:
corroded
BritanniaWater:fresh
BritanniaWater:
corroded
Iro
n (
mg/L
)
HOCl NH₂Cl
0.0
0.3
0.6
0.9
1.2
MannheimWater:fresh
MannheimWater:
corroded
BritanniaWater:fresh
BritanniaWater:
corroded
Copper
(mg/L
)
HOCl NH₂Cl
0.0
0.2
0.4
0.6
MannheimWater:fresh
MannheimWater:
corroded
BritanniaWater:fresh
BritanniaWater:
corroded
Lead (
mg/L
)
HOCl NH₂Cl
91
For either fresh or corroded metal coupons, generally, there were greater amounts of
metals released into Britannia Water than Mannheim Water. A paired student’s t test at a
confidence level of 95% was applied to all of the metal concentration values in Mannheim Water
and Britannia Water, and its p-value was 0.03, confirming that Britannia Water was statistically
more corrosive to metal coupons than Mannheim Water, likely due to the lower pH and
alkalinity values of Britannia Water.
For iron coupons, as shown in Figure 4-8, the impacts of HOCl and NH2Cl on the
released iron concentrations were inconsistent regarding each type of water and coupon ages. For
fresh iron coupons, NH2Cl increased the released iron concentrations in Mannheim Water but
exhibited no significant difference relative to HOCl in Britannia Water. For corroded iron
coupons, NH2Cl tended to be less aggressive in Mannheim Water but more aggressive in
Britannia Water. It is interesting that Cantor et al. (2003) and LeChevallier et al. (1990b) have
reported an increased iron corrosion rate under chlorination than under chloramination. Perhaps
they employed different water matrix and metal age, thereby leading to a different conclusion
about the impacts of disinfectant type on iron corrosion.
For fresh copper coupons, a higher concentration of copper was released under
chloramination than chlorination in both water matrices. This observation was different from the
trend reported in several studies, in which HOCl was more aggressive towards copper than
chloramine due to its higher oxidation potential (Boulay and Edwards, 2001; Boyd et al., 2008;
Cantor et al., 2003; Rahman et al., 2007). However, few studies have been conducted to compare
the impacts of disinfectants on copper release for new copper pipes. In chloraminated copper
plumbing, ammonia or ammonium ions, a byproduct from chloramine decay, tends to form
stable complexes with copper ions [Cu(I) and Cu(II)] and thus increase the solubility of copper
corrosion solids (Boyd et al., 2008; Schock and Lytle, 1995). Furthermore, the stability of
dissolved Cu(I)-ammonia complexes may hasten corrosion processes by reactions such as
[Cu(NH3)4]2+
+ Cu(s) 2[Cu(NH3)2]+ (Cotton and Wilkinson, 1988), especially in relatively
new copper pipes where a large area of pure copper is exposed to an oxidizing environment. For
corroded copper coupons, as shown in Figure 4-8, copper concentrations resulting from HOCl
exposure were approximately 2 fold higher than those under NH2Cl conditions in both water
matrices. This observation is consistent with the results of most studies in that HOCl is more
aggressive to copper corrosion than NH2Cl.
92
For both fresh and corroded lead coupons in Mannheim Water, as shown in Figure 4-8,
the released lead concentrations after 24 hours of chlorination (0.5 mg/L and 0.07 mg/L for fresh
and corroded coupons, respectively) were significantly higher than those under chloramination
(0.2 mg/L and 0.03 mg/L for fresh and corroded coupons, respectively). In Britannia Water,
however, there was no significant difference in lead concentrations between HOCl and NH2Cl
for fresh lead coupons after 24 hours, but NH2Cl tended to increase lead dissolution for corroded
coupons. Normally, the higher oxidation potential of HOCl promotes the formation of less
soluble lead dioxide (PbO2), whereas divalent lead solids [e.g., Pb3(CO3)2(OH)2] will form under
chloramination. As such, an increased lead level generally has been observed in chloraminated
water (Boyd et al., 2009; Edwards and Dudi, 2004). However, the transition kinetics of Pb(II) to
Pb(IV) depends on initial chlorine concentration, hydrodynamics, and water quality parameters
such as alkalinity and NOM. Only a dense surface layer of PbO2 could dramatically decrease
lead release (Boyd et al., 2008). Therefore, the observed increase in lead concentrations under
HOCl than those exposed to NH2Cl for fresh coupons in two water matrices were possibly
because these coupons had undergone oxidation by HOCl only for 24 hours and PbO2 was not
readily formed during the short-term reaction. Instead, Pb(II)-containing solids were most likely
the controlling solid phase for lead solubility, as was evidenced by the XRD results (Section
4.3.1). For corroded lead coupons in Mannheim Water, due to relatively high levels of alkalinity
and NOM in Mannheim Water, the transformation of Pb(II) to Pb(IV) was likely inhibited.
Therefore, the high HOCl oxidation potential also induced more lead to be released in the
chlorinated Mannheim Water. In contrast, for corroded coupons in Britannia Water, PbO2 was
likely the dominant solid phase under HOCl, and its low solubility determined that the lead
concentrations after a further 24 hour exposure was significantly reduced compared with that
under NH2Cl. Overall, the inconsistent effects of HOCl and NH2Cl on lead corrosion in two
water matrices indicate that water quality and metal age may confound the impacts of
disinfectant type, thereby affecting the compositions of solid controlling phase and the
subsequent dissolution of lead into the water (Edwards and Dudi, 2004; Kim and Herrera, 2010).
4.4 Summary
This chapter summarizes the metal species and concentration data that were obtained
during the experiments described in Chapter 3, presented separately here due to the quantity of
data. In this chapter, the effects of disinfectant type, metal age, water quality and phosphate-
93
based inhibitors on metal corrosion were evaluated. The test coupons of ductile iron, copper and
lead, representing either water main or premise plumbing, were investigated. The following
conclusions can be drawn based on the experimental results:
a. The observations concerning the impacts of orthophosphate on released iron, copper and
lead concentrations were generally consistent with those of other researchers.
Orthophosphate had no impacts on the released iron levels, regardless of metal surface
age, water quality, or disinfectant type. However, orthophosphate significantly reduced
the released copper concentrations, in particular for fresh copper surfaces and short term
exposures. Lead release was significantly reduced in the presence of orthophosphate,
irrespective of the age of the metal surface or water quality involved.
b. Results of XPS surface analysis suggest that the effectiveness of orthophosphate on
copper and lead corrosion control was due to facilitated precipitation of calcium by
orthophosphate to form calcium phosphate, which is consistent with previous studies.
The role of CaCO3 in determining the abundance of oxidized iron on the iron surface was
newly identified. Surface characterization of corroded iron coupons demonstrates that
CaCO3 precipitation decreased the amount of iron oxides formed on the iron surface.
c. In agreement with other research, high levels metal ions of iron, copper and lead were
released in waters with low pH and alkalinity.
d. The experimental design uniquely enabled the direct comparison of released metal
concentrations under chlorination and chloramination. However, the impacts of HOCl
and NH2Cl on the released iron, copper and lead concentrations for both fresh and pre-
corroded metal materials in the two tested water matrices were inconclusive. Only for
pre-corroded copper coupons was HOCl consistently more aggressive than NH2Cl in both
Mannheim and Britannia water.
In addition, since all of the data in this chapter were collected at the same time as those in
Chapter 3, they can also help provide some explanations for the trends in chlorine decay and
HAA formation that were described in Chapter 3. For example, for fresh copper coupons in
Mannheim water, phosphate-based corrosion inhibitors significantly sequestered copper release
compared with copper coupons in the absence of corrosion inhibitors. Due to the protection from
these corrosion inhibitors, the copper surface was also less readily available to react with
94
chlorine, leading to decreased free chlorine degradation rates in comparison with copper coupons
in the absence of corrosion inhibitors.
4.5 References
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Edition, Washington D C, USA.
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corrosion by-product release in soft water. Water Research, 35(3), 683-690.
Boyd, G. R., Dewis, K. M., Korshin, G. V., Reiber, S. H., Schock, M. R., Sandvig, A. M., and
Giani, R. (2008) Effects of changing disinfectants on lead and copper release. Journal
American Water Works Association, 100(11), 75-87.
Boyd, G. R., Dewis, K. M., Korshin, G. V., Reiber, S. H., Schock, M. R., Sandvig, A. M., and
Giani, R. (2009) Effects of Changing Disinfectants on Lead and Copper Release (vol 100,
pg 75, 2008). Journal American Water Works Association, 101(2), 10-10.
Brereton, J. A., and Mavinic, D. S. (2002) Field and material-specific simulated distribution
system testing as aids to understanding trihalomethane formation in distribution systems.
Canadian Journal of Civil Engineering, 29(1), 17-26.
Cantor, A. F., Denig-Chakroff, D., Vela, R. R., Oleinik, M. G., and Lynch, D. L. (2000) Use of
polyphosphate in corrosion control. Journal American Water Works Association, 92(2),
95-102.
Cantor, A. F., Park, J. K., and Vaiyavatjamai, P. (2003) Effect of chlorine on corrosion in
drinking water systems. Journal American Water Works Association, 95(5), 112-123.
Cotton, F. A., and Wilkinson, G. (1988) Advanced Inorganic Chemistry. Fifth Edition. John
Wiley & Sons, Inc., New York.
Dartmann, J., Alex, T., Dorsch, T., Schevalje, E., and Johannsen, K. (2004) Influence of
decarbonisation and phosphate dosage on copper corrosion in drinking water systems.
Acta Hydrochimica Et Hydrobiologica, 32(1), 25-32.
95
Demora, S. J., and Harrison, R. M. (1984) Lead in tap water - contamination and chemistry.
Chemistry in Britain, 20(10), 900-904.
Edwards, M., and Dudi, A. (2004) Role of chlorine and chloramine in corrosion of lead-bearing
plumbing materials. Journal American Water Works Association, 96(10), 69-81.
Edwards, M., Hidmi, L., and Gladwell, D. (2002) Phosphate inhibition of soluble copper
corrosion by-product release. Corrosion Science, 44(5), 1057-1071.
Edwards, M., and McNeill, L. S. (2002) Effect of phosphate inhibitors on lead release from
pipes. Journal American Water Works Association, 94(1), 79-90.
Edwards, M., Schock, M. R., and Meyer, T. E. (1996) Alkalinity, pH, and copper corrosion by-
product release. American Water Works Association Journal, 88(3), 81-94.
Hach (2007) Chloramine (Mono) - Indophenol Method 10171 (DOC316.53.01015), edition 6.
Hach Company, Loveland, Colorado.
Health Canada (2009) Guidance on Controlling Corrosion in Drinking Water Distribution
Systems (Catalogue No. H128-1/09-595E). Water, Air and Climate Change Bureau,
Healthy Environments and Consumer Safety Branch, Health Canada, Ottawa, Ontario.
Hsu, P. H. (1982) Crystallization of Iron(Iii) Phosphate at Room-Temperature. Soil Science
Society of America Journal, 46(5), 928-932.
Kim, E. J., and Herrera, J. E. (2010) Characteristics of Lead Corrosion Scales Formed during
Drinking Water Distribution and Their Potential Influence on the Release of Lead and
Other Contaminants. Environmental Science & Technology, 44(16), 6054-6061.
LeChevallier, M. W., Lowry, C. D., and Lee, R. G. (1990) Disinfecting Biofilms in a Model
Distribution System. American Water Works Association Journal, 82(7), 87-99.
Leroy, P. (1993) Lead in drinking water: Origins, solubility, treatment. Aqua (Oxford), 42(4),
233-238.
96
MacQuarrie, D. M., Mavinic, D. S., and Neden, D. G. (1997) Greater Vancouver Water District
drinking water corrosion inhibitor testing. Canadian Journal of Civil Engineering, 24(1),
34-52.
Maddison, L. A., Gagnon, G. A., and Eisnor, J. D. (2001) Corrosion control strategies for the
Halifax regional distribution system. Canadian Journal of Civil Engineering, 28(2), 305-
313.
McNeill, L. S., and Edwards, M. (2000) Phosphate inhibitors and red water in stagnant iron
pipes. Journal of Environmental Engineering-Asce, 126(12), 1096-1102.
McNeill, L. S., and Edwards, M. (2001) Iron pipe corrosion in distribution systems. American
Water Works Association Journal, 93(7), 88-100.
McNeill, L. S., and Edwards, M. (2004) Importance of Pb and Cu particulate species for
corrosion control. Journal of Environmental Engineering-Asce, 130(2), 136-144.
Moriarty, B. E. (1990) Surface Studies of Corrosion-Inhibitors in Cooling Water-Systems.
Materials Performance, 29(1), 45-48.
NACE International (2009) NACE International glossary of corrosion-related terms. National
Association of Corrosion Engineers International, Houston, TX.
Persson, P., Nilsson, N., and Sjoberg, S. (1996) Structure and bonding of orthophosphate ions at
the iron oxide aqueous interface. Journal of Colloid and Interface Science, 177(1), 263-
275.
Rahman, S., McDonald, B. C., and Gagnon, G. A. (2007) Impact of secondary disinfectants on
copper corrosion under stagnation conditions. Journal of Environmental Engineering-
Asce, 133(2), 180-185.
Sarin, P., Clement, J. A., Snoeyink, V. L., and Kriven, W. W. (2003) Iron release from corroded
unlined cast-iron pipe. Journal American Water Works Association, 95(11), 85-96.
Schock, M. R. (1980) Response of lead solubility to dissolved carbonate in drinking water.
Journal American Water Works Association, 72(12), 695-704.
97
Schock, M. R. (1989) Understanding corrosion control strategies for lead. Journal American
Water Works Association, 81(7), 88-100.
Schock, M. R. (1990) Causes of temporal variability of lead in domestic plumbing systems.
Environmental Monitoring and Assessment, 15(1), 59-82.
Schock, M. R., and Lytle, D. A. (1995) Effect of pH, DIC, Orthophosphate and Sulfate on
Drinking Water Cuprosolvency. U.S. Environmental Protection Agency, Water Supply
and Water Resources Division, Cincinnati, Ohio 45268, EPA/600/R-95/085.
Schock, M. R., and Sandvig, A. M. (2009) Long-term effects of orthophosphate treatment on
copper concentration. Journal American Water Works Association, 101(7), 71-82.
Scholze R.J.; Pontow K.A.; Kanchlbhatla G. and Ray B.T. (1994) Using the CERL pipe loops
system (PLS) to evaluate corrosion inhibitors that can reduce lead in drinking water
(FEAP-TR-EP-94/04). U.S. Army Construction Engineering Research Laboratories,
Champaign, IL.
Singley, J. E. (1994) Electrochemical nature of lead contamination. Journal American Water
Works Association, 86(7), 91-96.
Stone, E. D. (2008) Effects of Orthophosphate Corrosion Inhibitor in Blended Water Quality
Environments. PhD Dissertation, University of Central Florida.
Summers, R. S., Hooper, S. M., Shukairy, H. M., Solarik, G., and Owen, D. (1996) Assessing the
DBP yield: Uniform formation conditions. Journal American Water Works Association,
88(6), 80-93.
Tang, Z., Hong, S., Xiao, W., and Taylor, J. (2006) Characteristics of iron corrosion scales
established under blending of ground, surface, and saline waters and their. Corrosion
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Vasquez, F. A., Heaviside, R., Tang, Z. J., and Taylor, J. S. (2006) Effect of free chlorine and
chloramines on lead release in a distribution system. Journal American Water Works
Association, 98(2), 144-154.
98
Williams, S. M. (1990) The use of sodium-silicate and sodium polyphosphate to control water-
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Xiao, W. Z., Hong, S. K., Tang, Z. J., and Taylor, J. S. (2007) Effects of blending on total copper
release in distribution systems. Journal American Water Works Association, 99(1), 78-88.
99
Catalytic Impacts of Copper Corrosion Products on 5
Chlorine Decay and HAA Formation in Simulated
Distribution Systems
The results in this chapter have been published as:
Zhang, H., Andrews, S.A. (2012) Catalysis of Copper Corrosion Products on Chlorine Decay
and HAA Formation in Simulated Distribution Systems. Water Research, 46 (8), 2665-2673.
Results from this chapter focus on the research gap “What are the impacts of metal
corrosion products on the fate of HAA in distribution systems?”
100
Abstract
This study investigated the effect of copper corrosion products, including Cu(II), Cu2O,
CuO and Cu2(OH)2CO3, on chlorine degradation, HAA formation, and HAA speciation under
controlled experimental conditions. Chlorine decay and HAA formation were significantly
enhanced in the presence of copper with the extent of copper catalysis being affected by the
solution pH and the concentration of copper corrosion products. Accelerated chlorine decay and
increased HAA formation were observed at pH 8.6 in the presence of 1.0 mg/L Cu(II) compared
with that observed at pH 6.6 and pH 7.6. Further investigation of chlorine decay in the presence
of both Suwannee River NOM and Cu(II) indicated that an increased reactivity of NOM via
interacting with dissolved and/or solid surface-associated Cu(II), rather than chlorine auto-
decomposition, was a primary reason for the observed rapid chlorine decay. Copper corrosion
solids [Cu2O, CuO, Cu2(OH)2CO3] exhibited catalytic effects on both chlorine decay and HAA
formation. Contrary to the results observed when in the absence of copper corrosion products,
DCAA formation was consistently predominant over other HAA species in the presence of
copper corrosion products, especially at neutral and high pH. This study improves the
understanding for water utilities and households regarding chlorine residuals and HAA
concentrations in distribution systems, in particular once the water reaches domestic plumbing
where copper is widely used.
Keywords: Copper; Corrosion; Catalysis; Free chlorine; Haloacetic acids
101
5.1 Introduction
To maintain the microbial stability of distributed water, free chlorine is used as a
secondary disinfectant in distribution systems. As chlorine may experience temporal and spatial
degradation due to chemical and biological consumption that occurs in the bulk water and on the
pipe wall, at least 0.2 mg/L of free chlorine residual should be maintained to reduce the possible
occurrence of biological regrowth (USEPA, 1989).
Haloacetic acids (HAAs) are formed primarily by the chlorination of natural organic
matter (NOM). In its simplest terms, NOM can be classified as either humic substances or non-
humic compounds. However, humic substances are a mixture of many molecules, and the major
functional groups which contribute to surface charge and reactivity of humic substances include
carboxyl groups, some phenolic groups, alcohol groups, methoxyl groups, ketones, and
aldehydes (Reckhow et al., 1990). Humic substances in the environment are capable of
interacting with metal ions to form soluble complexes, colloidal substances and/or insoluble
substances (Stevenson, 1994). Alvarez-Puebla et al. (2004a, 2004b) have investigated possible
retention mechanisms of Cu(II), Co(II) and Ni(II) on humic substances and found surface
complexation and electrostatic retention played key roles in the retention of these metals on
humic substances, and Cu(II) had a higher affinity for humic substances compared with Co(II)
and Ni(II). Furthermore, the humic substances studied displayed a great selectivity for different
Cu(II) species, and the selectivity depended on pH.
Since the 1970s, considerable effort has been made to understand HAA formation
mechanisms by using NOM or well-defined model compound precursors (Kanokkantapong et
al., 2006; Morris, 1975; Reckhow et al., 1990). Reaction mechanisms involved in HAA
formation have been found to generally include oxidation, substitution, addition, and hydrolysis
(Morris, 1975). The reaction rates and HAA speciation are dependent on temperature, pH,
chlorine dose and the nature of the organic compounds that contribute to NOM. Although there
are nine HAA species, monochloroacetic acid (MCAA), dichloroacetic acid (DCAA),
trichloroacetic acid (TCAA), monobromoacetic acid (MBAA) and dibromoacetic acid (DBAA)
are commonly detected in drinking water, and DCAA and TCAA are normally dominant over
other species. Due to the toxicological properties of HAAs, the Stage 2
102
Disinfectant/Disinfection Byproduct Rule (D/DBPR) has regulated a Maximum Contaminant
Level (MCL) of HAA5 (MCAA, DCAA, TCAA, MBAA, and DBAA) at 0 μg L (USEPA,
2006). The Guidelines for Canadian Drinking Water Quality have established a Maximum
Acceptable Concentration (MAC) for HAA5 in drinking water at 80 μg L, based on a locational
running annual average of a minimum of quarterly samples taken in the distribution system
(Health Canada, 2008).
Copper corrosion occurs as water containing a disinfectant residual travels through
copper pipes. Copper ions [Cu(II)], cuprite [Cu2O], tenorite [CuO], cupric hydroxide [Cu(OH)2],
and malachite [Cu2CO3(OH)2] are dominant corrosion products of copper in drinking water
systems (Xiao et al., 2007). Copper has no acute toxicity to humans except at high dose (>15
mg/day). An aesthetic objective of ≤1.0 mg/L has been established for copper in drinking water
(Health Canada, 1992).
The catalytic potential of copper in many reaction processes has been widely investigated
(Onuchukwu, 1994; Paidar et al., 1999; Pintar et al., 1997). Other studies have also demonstrated
copper catalysis on free chlorine and monochloramine degradation and on THM formation
(Blatchley et al., 2003; Fu et al., 2009a; Fu et al., 2009b; Li et al., 2007; Li et al., 2008). The
mechanism of Cu-catalyzed THM formation proposed by Blatchley et al. (2003) was that copper
could complex with THM precursor compounds and enhance the oxidative decarboxylation and
enolization of the keto-groups.
Since HAA formation has been found to be concurrent with THM formation, with both
DBP classes having similar organic precursors and some similarities in their formation
mechanisms (Morris, 1975; Reckhow and Singer, 1985), and elevated HAA formation in the
presence of high levels of copper ion was observed in Chapter 3, it is hypothesized that copper
will play a similar catalytic role in HAA formation. Li et al. (2008) have reported elevated
formation of HAAs in copper pipes. However, they did not examine in detail the catalytic
potential of Cu(II) and other corrosion products to affect chlorine degradation and HAA
formation. Furthermore, a regulatory guideline for HAA5 has recently been added to the
“Guidelines for Canadian Drinking Water Quality” (Health Canada, 2008). This information,
therefore, has increased the awareness of water utilities and households to minimize HAAs in
their water systems. As such, the objective of this study was to investigate the effect of Cu(II)
103
and its solid corrosion products on chlorine degradation and HAA formation, including HAA
speciation, by applying different concentrations of copper corrosion products and under different
pH conditions. Since copper is widely used in domestic plumbing systems, understanding the
roles of copper and its corrosion products on the fate of HAAs will benefit utilities and
households to minimize HAAs in their systems by adopting preventive strategies for their
control.
5.2 Materials and Methods
Reagents and Materials 5.2.1
All chemicals used in this study were ACS grade or higher. The chlorine dosing solution
(approximately 3500 mg/L as Cl2) was prepared by diluting a concentrated solution of sodium
hypochlorite (NaOCl, 6%, VWR) in Milli-Q water. The target chlorine concentration applied in
the tests was achieved by spiking 3.5 mL chlorine dosing solution in 1 L test water. An initial
free chlorine concentration of 10 mg/L as Cl2 was applied to achieve a detectable 24-h residual of
0.2 mg/L as Cl2, which was found to be necessary especially when the effects of Cu2O were
investigated. Phosphate buffer (H2PO4-/HPO4
2-) and borate buffer (H3BO3/NaOH) solutions were
prepared and added to test solutions at 1 mM concentration to control reaction solution pH at
desired levels. Sulfuric acid (H2SO4, 50%) and sodium hydroxide (NaOH, 20%) were also used
for pH adjustment. L-Ascorbic acid (99.0%, Sigma Aldrich) at 200 mg/L was applied to quench
chlorine residuals before HAAs extraction. Over the reaction period, the pH was maintained at
their specified target values (± 0.1).
Unchlorinated post-filtration water was collected from the Mannheim Water Treatment
Plant (MWTP), Ontario, for testing. Water quality parameters are listed in Table 5-1. In the tests
to determine reaction pathways of Cu(II) catalysis on chlorine decay, synthetic water with
similar inorganic chemical composition to the water from the MWTP was prepared. Suwannee
River NOM (International Humic Substances Society) was dosed into synthetic water at the
concentration of 2.4 mg/L and 4.7 mg/L.
104
Table 5-1Water quality parameters for post-filtration water from MWTP
Parameters Parameter values
pH 7.5 ± 0.2
UV254 (cm-1
) 0.045 ± 0.015
TOC (mg/L) 3.9 ± 0.4
SUVA (L/mg·cm-1
) 0.016 ± 0.004
Bromide (μg L) 65.0 ± 15.5
Chloride (mg/L) 84.5 ± 2.5
Sulfate (mg/L) 35.0 ± 2.0
Cl-:SO4
2- ratio 2.4 ± 0.1
Experimental Procedures 5.2.2
All of the experiments were performed in 1 L amber bottles with PTFE-lined caps and at
room temperature (21±2 ºC). The reaction bottles were made chlorine demand free before use.
All of the tests also included control samples, which were prepared in the same way as for test
samples but without the addition of copper corrosion products.
Copper ions were spiked in the form of copper nitrate [Cu(NO3)2] at concentrations of 0,
0.2, 1.0 and 2.0 mg/L of copper to be compatible with the common range of copper
concentrations in distribution systems of 20-2020 µg/L (Health Canada, 1992). Cuprous oxide
(Cu2O), cupric oxide (CuO) and malachite [Cu2CO3(OH)2] were selected as the solid forms of
copper corrosion products of interest due to their prevalence in copper pipes. They were added
into test solutions in the form of powders and mixed with test water thoroughly using a stir bar.
The concentrations investigated in this study were 0, 0.2 and 1.0 g/L, which were comparable to
the dosages applied by Li et al. (2007). After each prescribed reaction time (between 2 and 120
hours), water samples dosed with corrosion products were taken from the reaction bottles and
filtered through 0.2 µm Nylaflo® Nylon membrane filter paper (Pall Corporation) to separate the
solid copper corrosion products from the aqueous solution. Dissolved copper in the filtrate was
measured by flame atomic absorption spectrometry (FAAS). To account for any influence from
filter paper and filtration performance on the concentration of chlorine residual and HAA9,
control samples were also filtered under similar conditions. When the effects of pH on chlorine
decay and HAA formation were investigated (Sections 5.3.1.1 and 5.3.2.1), the solution pH was
controlled with 1 mM phosphate or borate buffer at pH 6.6, 7.6 or 8.6 to represent a possible pH
range in distribution systems. In the tests to investigate the effects of copper concentration
(Sections 5.3.1.3 and 5.3.2.2) and solid copper corrosion products (Sections 5.3.1.4 and 5.3.2.3),
105
and the tests to determine reaction pathways of Cu(II) catalysis on chlorine decay (Section
5.3.1.2), the pH of both natural and synthetic water was controlled at 8.3 ± 0.1 by the addition of
1 mM borate buffer.
In addition, to evaluate the impacts of NOM on copper solubility, copper ions were
spiked into Milli-Q water and MWTP water in the form of Cu(NO3)2 at concentrations of 0, 0.2,
1.0 and 2.0 mg/L of copper, and pH was controlled at 8.3 with 1 mM borate buffer. The copper-
spiked water was then filtered through 0.2 µm Nylaflo® Nylon membrane filter paper, and
copper concentrations in the filtrate were measured by FAAS.
In this study, free chorine decay rates (Section 5.3.1) were estimated by fitting a pseudo-
first-order decay equation to free chlorine data using Microsoft Excel. HAA formation rates
(Section 5.3.2) were determined by fitting HAA concentrations versus reaction time with a
logarithmic function in Microsoft Excel. A one-factor analysis of variance (ANOVA) test at a
confidence level of 95% was applied to determine whether each treatment factor, including pH
and Cu(II) concentrations, had significant impacts on free chlorine decay and HAA formation
(Montgomery, 2000).
Nine species of haloacetic acids (HAA9) were analyzed according to EPA Standard
Method 6251 B with 2,3,5,6-tetrafluorobenzoic acid as a surrogate standard (APHA, AWWA,
WEF, 2005). Measurement of total organic carbon (TOC) concentration was undertaken using a
Model 1030 TOC analyzer (OI Analytical, USA). UV254 was measured using a CE 3055
Reflectance Spectrophotometer (Cecil Instruments, England) at 254 nm. Measurements of free
chlorine were performed using a Hach DR2700 Spectrophotometer (Hach Company, USA) at
530 nm. Soluble copper concentrations in the filtered water samples were analyzed using a
Varian SpectrAA.20 Flame Atomic Absorption Spectrometer (Agilent Technologies, USA) after
being acidified to pH <2.
5.3 Results and Discussion
The concentrations of free chlorine and HAA9 were monitored by periodically
withdrawing and analyzing samples from the test bottles. Therefore, it was possible to compare
free chlorine decay rates and HAA9 formation kinetics among samples with and without spikes
106
of Cu(II) and/or copper corrosion products and evaluate their catalytic potential on free chlorine
degradation and HAA formation.
Chlorine Decay 5.3.1
5.3.1.1 Effects of pH
In the absence of Cu(II), there was no statistical difference in chlorine decay among the
examined three pH values (p-value >0.05). Pseudo-first-order decay rates were estimated to be
0.0018, 0.0022, and 0.0022 h-1
for pH 6.6, pH 7.6 and pH 8.6, respectively. These results are in
agreement with previous findings (Powell et al., 2000). However, when 1 mg/L Cu(II) was spiked,
as shown in Figure 5-1, free chlorine degraded more rapidly at higher pH. Compared with chlorine
decay rates in the absence of Cu(II), no significant effect was observed at pH 6.6 (p-value >0.05),
but chlorine degradation was significantly accelerated at pH 7.6 and 8.6 (p values of 0.02 and
0.003, respectively). Pseudo first-order decay rates in the presence of Cu(II) at pH 7.6 and 8.6 were
0.0033 and 0.0061 h-1
, respectively.
Figure 5-1 Chlorine degradation at different pH values in the absence and presence of 1 mg/L Cu
(II) for MWTP water; initial HOCl =10 mg/L; triplicate
Both MINEQL+ version 4.5 (Environmental Research Software, Hallowell, Maine, USA)
and some simple filtration tests were used to determine if the spiked copper remained dissolved
during these tests or pH manipulations caused some solid species to form. At pH 6.6 and 7.6,
malachite [Cu2(OH)2CO3] was estimated by MINEQL+ to be the primary form of the solids, and
Cu(OH)2 was the dominant solid at pH 8.6. The concentrations of these solid copper species were
estimated to increase with increasing pH, with 76%, 93% and 100% of the total 1 mg/L spiked
0.0
0.2
0.4
0.6
0.8
1.0
0 20 40 60 80 100 120
C/C
0
Time (hours)
pH 6.6-with Cu
pH 7.6-with Cu
pH 8.6-with Cu
107
copper being transformed into solids at pH 6.6, pH 7.6 and pH 8.6, respectively. These results
were confirmed by filtering and measuring dissolved copper concentrations in the filtrate of
samples that had been prepared in Milli-Q water. However, even at pH 8.6, measured dissolved
copper concentrations in the filtered MWTP water accounted for 93% of the spiked total 1 mg/L
copper, leaving only 7% of copper possibly being present in solid form. This significant increase in
copper solubility in natural water, relative to the results by MINEQL simulation for Milli-Q water
support reports that the presence of NOM in natural water facilitates the formation of Cu-NOM
complexes due to its strong binding affinities and thus hindering the formation of solid copper
species (Korshin et al., 1996). Regardless, although copper had been added in its highly soluble
Cu(NO3)2 form, it did not always remain in dissolved form and many samples had both dissolved
and solid copper species present. As such, at this point in the experiments, faster degradation of
chlorine at higher pH in the presence of Cu(II) as observed herein was hypothesized to have been
due to an accelerated reactivity of organic precursors that had interacted with dissolved and/or
solid copper species at high pH (Blatchley et al., 2003; Fu et al., 2009b) and/or the decomposition
of chlorine (as OCl-) to O2 and Cl
- by Cu(II) in alkaline solutions (Gray et al., 1977). Experiments
described in the following sections were performed to investigate these possibilities.
5.3.1.2 Effects of NOM Concentration
The relative contribution of auto-decomposition and consumption (through reacting with
organic matter) to the observed rapid chlorine decay in the presence of Cu(II) at high pH was
examined using synthetic water with similar inorganic chemical composition to the test water
described in Section 5.3.1.1. Chlorine decay kinetics were compared between solutions with and
without the addition of 1 mg/L Cu(II) and Suwannee River NOM.
As shown in Figure 5-2, in the absence of Suwannee River NOM, no statistical difference
in chlorine decay was observed when comparing results obtained with synthetic water and
synthetic water spiked with Cu(II) (p-value >0.05). This indicates that chlorine auto-
decomposition was not affected by copper catalysis within the 25-h test period of this study.
When 2.4 mg/L Suwannee River NOM was spiked, chlorine degraded significantly faster
(k=0.012 h-1
) with a p-value of 0.007 compared with water in the absence of NOM (k=0.001 h-1
)
as a result of the consumption of chlorine by NOM. When 1.0 mg/L Cu(II) was added to NOM-
spiked water, the chlorine decay rate further increased to 0.048 h-1
, significantly higher than that
108
without the addition of Cu(II) (p-value of 0.001). Figure 5-2 also shows that as the NOM
concentration was doubled in the absence and presence of 1.0 mg/L Cu(II), the chlorine decay
rates increased to 0.055 h-1
and 0.584 h-1
, respectively. Almost no residual was detected after 6
hours for water spiked with 4.7 mg/L NOM and 1.0 mg/L Cu(II). These results suggest that the
observed rapid chlorine degradation was primarily attributed to interactions of NOM with Cu(II)
(dissolved and/or solid phases) and subsequent increases in NOM reactivity with free chlorine.
Figure 5-2 Chlorine decay for synthetic water in the presence and absence of 1 mg/L Cu(II) and
NOM; initial HOCl =4.2 mg/L; pH 8.3; duplicate
5.3.1.3 Effects of Copper Concentration
Cu(NO3)2 were spiked into MWTP water at concentrations of 0 mg/L, 0.2 mg/L, 1.0
mg/L, and 2.0 mg/L of copper. Approximately 10 mg/L chlorine was spiked, and the chlorine
residual was monitored for the following 100 hours. Figure 5-3 compares the observed chlorine
degradation kinetics with respect to the different spiked Cu(II) concentrations. The chlorine
degradation rates significantly increased with an increase in the spiked Cu(II) concentrations
from 0 to 1.0 mg/L (p-value of 1.810-10
), but no additional difference in chlorine degradation
rate was observed for water dosed between 1.0 mg/L and 2.0 mg/L Cu(II) (p-value >0.05). The
increased chlorine degradation in the presence of Cu(II) may be explained by copper’s potential
to complex with organic compounds and enhance DBP formation by accelerating the
decarboxylation and enolization steps involved in DBP formation pathways (Blatchley et al.,
2003). Consequently, more chlorine may be involved in electrophilic addition and substitution
reactions with organic precursors. Since chlorine decay and DBP formation are not independent
0.0
0.3
0.6
0.9
1.2
0 5 10 15 20 25
C/C
0
Time (hours)
Synthetic water
Synthetic water+Cu(II)
Synthetic water+2.4 mg/LNOM
Synthetic water+2.4 mg/LNOM+Cu(II)
Synthetic water +4.7mg/L NOM
Synthetic water+4.7 mg/LNOM+Cu(II)
k=0.012 h-1
k=0.048 h-1
k=0.055 h-1
k=0.584 h-1
k=0.001 h-1
109
phenomena, further discussion of the effects of Cu(II) concentration on HAA formation is
provided in Section 5.3.2.2.
Figure 5-3 Chlorine degradation in the presence of different concentrations of Cu(II) for MWTP
water; initial HOCl =10 mg/L; triplicate; pH 8.3
5.3.1.4 Effects of Solid Copper Corrosion Products
CuO, Cu2O and Cu2(OH)2CO3, all commonly identified copper corrosion products, were
studied to identify their effects on chlorine degradation. Compared with control samples,
chlorine degraded more quickly in the presence of these corrosion solids. The observed sequence
of chlorine degradation in the presence of copper corrosion solids was 1.0 g/L Cu2O > 0.2 g/L
Cu2O > 1.0 g/L malachite > 0.2 g/L malachite 1.0 g/L CuO > control (Figure 5-4). The
concentrations of Cu(II) that are shown for solutions containing these solid corrosion products
were the equilibrated concentrations of Cu(II) released from the solids. Released Cu(II)
concentrations that were measured after 20 hours of reaction were at levels of approximately 0.8
mg/L for solutions of both 1.0 g/L Cu2O and 0.2 g/L Cu2O, and approximately 0.2 mg/L for
solutions of 1.0 g/L and 0.2 g/L malachite. These results indicate that the dissolved Cu(II)
reached an approximate equilibrium with the solids during the tests and that the equilibrium
Cu(II) concentrations measured agree with published solubility data for these compounds (Li et
al., 2008; Merkel et al., 2002).
0.0
0.2
0.4
0.6
0.8
1.0
0 20 40 60 80 100 120
C/C
0
Time (hours)
0 mg/L
0.2 mg/L
1.0 mg/L
2.0 mgL
110
Figure 5-4 Pseudo-first-order decay rates of free chlorine for MWTP water containing dissolved
Cu(II) and solid copper corrosion products
Li et al. (2007, 2008) have reported that the effects of CuO and Cu2O on the degradation
of chlorine were actually due to the effect of Cu(II) released from the oxides. To determine if the
released Cu(II) from copper corrosion solids also played a significant role in the observed
difference in chlorine decay kinetics among these solid corrosion products in the present study,
the pseudo-first-order decay constants for free chlorine were compared with those for free
chlorine in contact with only dissolved Cu(II) in the filtered samples spiked with Cu(NO3)2
(Figure 5-4). In all cases, the decay constants were obtained by fitting a pseudo first-order decay
equation to the chlorine data once the released Cu(II) concentration reached equilibrium after
approximately 20 hours reaction time.
Chlorine decay rates were higher in the presence of both solid corrosion products (0.2-
1.0 g/L) and their associated dissolved Cu(II) (0.1-0.8 mg/L) than in the solutions spiked with
only Cu(NO3)2 (up to 2 mg/L of copper). For reactions involving only the dissolved form of
Cu(II), free chlorine decay rates increased approximately linearly with increases in the dissolved
Cu(II) concentrations from 0 mg/L to 1.0 mg/L (0.004 h-1
- 0.0117 h-1
) as shown in the inset plot
in Figure 5-4. In the presence of solid copper corrosion products, free chlorine decay rates were
significantly higher relative to those for solutions with similar aqueous dissolved Cu(II)
concentrations. For example, the decay constant for free chlorine in contact with 1.0 g/L CuO
was 0.0193 h-1
and the released Cu(II) concentration at equilibrium was 0.15 mg/L; whereas a
decay rate of only 0.0053 h-1
would be estimated from the curve at the same concentration of
0.0
0.1
0.2
0.3
0.0 0.2 0.4 0.6 0.8 1.0
Chlo
rine d
ecay c
onsta
nt (h
-1)
Measured dissolved Cu(II) (mg/L)
Cu(II) without solids 0.2 g/L Cu2O
1.0 g/L Cu2O 1.0 g/L CuO
0.2 g/L malachite 1.0 g/L malachite
0.000
0.005
0.010
0.015
0.020
0.0 0.2 0.4 0.6 0.8 1.0
111
dissolved Cu(II) in the absence of corrosion solids. In addition, however, increases in mass
concentrations of solids did not lead to the same magnitudes of increases in free chlorine
degradation rates. Therefore, the quantitative relationship between mass concentrations of copper
corrosion solids and free chlorine degradation needs further investigation. Regardless, the
observed significant increase in the decay constants for free chlorine in the presence of these
solid corrosion products indicates that not only the dissolved copper species, but also surface-
associated Cu(II) from these corrosion products, may be involved in the reactions and accelerate
free chlorine degradation.
These results are in contrast to those reported by Li et al. (2007, 2008) who discovered
that rapid chlorine decay in the presence of CuO and Cu2O was actually due to the effect of
Cu(II) released from the oxides. Therefore, differences in surface area of the solids may explain
some of the differences in the results from these two studies. For reactions mediated by surface-
bound metal, the reaction rate is affected by the mineral surface area and the density of sorbed
metal ions (Chun et al., 2005; Lee et al., 2008). Qualitatively, it follows that the total surface
area and hence the concentrations of reactive sites from 1.0 g/L Cu2O and malachite are larger
than for 0.2 g/L Cu2O and malachite, respectively. As a result, it is not surprising that free
chlorine degraded faster with 1.0 g/L Cu2O and malachite than with 0.2 g/L Cu2O and malachite.
However, due to the lack of information about surface area for the Cu2O, CuO and malachite
investigated in this study, normalization of the measured decay constants and subsequent more
quantitative comparisons of the reactivity of these solid corrosion products with free chlorine
cannot be made.
HAA Formation and Speciation 5.3.2
5.3.2.1 Effects of pH
The effect of pH on Cu(II) catalysis during HAA formation was investigated by
controlling the pH of MWTP water at 6.6, 7.6 and 8.6 and comparing HAA formation between
Cu(II)-spiked and unspiked water samples. Results indicate that copper catalysis during HAA
formation was highly dependent on pH, and the degree of enhancement of HAA formation by
Cu(II) became more apparent at higher pH values. After 100 hours of reaction time, as shown in
Figure 5-5, there was no statistical difference in HAA formation before and after the addition of
112
Cu(II) at pH 6.6 and 7.6 (p values >0.05). However, at pH 8.6, HAA formation significantly
increased with 1 mg/L Cu(II) as compared with the control (p-value of 0.006). Since
complexation of Cu(II) with organic matter is formed via Cu(II) interactions with carbonyl and
hydroxyl groups (Alvarez-Puebla et al., 2004b), and copper was present primarily in the form of
complexes with NOM in MWTP water at high pH (evidenced by 93% of copper being present in
the filtrate after passage through 0.2 µm filter paper), the observed effects of pH on HAA
formation in the presence of Cu(II) suggest that high pH promoted the electrostatic interactions
of dissolved and/or solid surface-associated Cu(II) with organic matter by favoring the ionization
of the aforementioned carbonyl and hydroxyl groups (Blatchley et al., 2003). Accelerated base-
catalyzed enolization and hydrolysis involved in HAA formation from a model precursor has
also been shown to contribute to elevated HAA formation in the presence of Cu(II) (Deborde and
von Gunten, 2008).
Figure 5-5 HAA9 at 100 hours at different pH values in the absence and presence of 1 mg/L
Cu(II) for MWTP water, initial HOCl =10 mg/L (error bars represent standard deviation of
triplicate tests)
5.3.2.2 Effects of Copper Concentration
The catalytic effect of Cu(II) on HAA formation under conditions of varying Cu(II)
concentrations is illustrated in Figure 5-6. HAA9 had similar formation kinetics in the first 30
hours for all of the water samples tested. When reactions proceeded further (t >70 hours), more
HAA compounds were produced for the water spiked with higher concentrations of Cu(II). There
was a significant increase in HAA9 formation with 1.0 mg/L Cu(II) after 100 hours of reaction
compared with HAA9 concentrations formed in control water (p-value of 3.210-6
). However, a
0
50
100
150
200
250
pH 6.6 pH 7.6 pH 8.6
HA
A9 (
µg
/L)
Cu(II) 0 mg/L Cu(II) 1.0 mg/L
113
further increase in the spiked copper concentration from 1.0 mg/L to 2.0 mg/L did not affect
HAA yields or formation rates significantly (p-value >0.05). Although the ability of Cu(II) to
complex with organic matter is essential for copper to act as a catalyst (Bansal et al., 2008), the
Cu(II) binding capacity of humic acid is limited, having been measured in the range of 48 to 160
mg/g of humic acid (Stevenson, 1985). The absence of a difference in HAA formation for the
spiked Cu(II) concentrations between 1.0 mg/L and 2.0 mg/L suggests that the functional groups
or binding sites of the organic molecules in the solution could have been saturated by 1.0 mg/L
of copper, regardless of its speciation. Since the reactivity of NOM with chlorine was enhanced
by the retention of dissolved and/or surface-bound Cu(II) on NOM (Sections 5.3.1.2, 5.3.1.3, and
5.3.1.4), it follows that the reactivity of NOM with chlorine would be increased with increasing
Cu(II) concentration until the active sites on NOM are completely occupied by Cu(II).
Figure 5-6 HAA9 formation in the presence of Cu(II) with varying concentrations at pH 8.3 for
MWTP water (error bars represent standard deviation of triplicate tests)
5.3.2.3 Effects of Solid Copper Corrosion Products
Copper corrosion solids were also observed to catalyze HAA formation at pH 8.3, the
extent of the effect being dependent on their species and mass concentrations. Compared with
the control sample, 1.0 g/L CuO had similar HAA formation, but 0.2 g/L Cu2O significantly
promoted HAA formation (p-value of 0.001) (Figure 5-7). No statistical difference in HAA
formation was observed for waters spiked with 0.2 g/L malachite and the control samples.
However, in the water spiked with 1.0 g/L of malachite, the HAA9 formation also significantly
increased compared with the control sample (p-value of 0.008).
0
30
60
90
120
150
0 20 40 60 80 100 120
HA
A9 (
µg
/L)
Time (hours)
0 mg/L
0.2 mg/L
1.0 mg/L
2.0 mg/L
114
Figure 5-7 HAA9 at a reaction time of 32 hours in the presence of copper corrosion solids at pH
8.3 for MWTP water (error bars represent standard deviation of triplicate tests)
Figure 5-8 enables a comparison of HAA formation rates for solutions spiked with
Cu(NO3)2 (dissolved Cu(II) without solids) and those spiked with solid copper corrosion
products. The concentrations of Cu(II) that are shown in Figure 5-8 for these solid corrosion
products were the measured equilibrated concentrations of Cu(II) released from the solids. As
shown in Figure 5-8, HAA formation rates for solutions spiked with the dissolved form of Cu(II)
increased with increases in the Cu(II) concentrations. The addition of copper corrosion solids
significantly increased HAA formation rates relative to those on the curve for solutions with
similar aqueous Cu(II) concentrations. This suggests that surface-associated Cu(II) from these
corrosion products may also participate in the reactions for HAA formation by interacting with
NOM and hence increasing its reactivity with free chlorine. As a result, free chlorine decay and
HAA formation were accelerated. As reported in Section 5.3.1.4, 0.2 g/L Cu2O and 1.0 g/L
malachite accelerated free chlorine degradation relative to 1.0 g/L CuO and 0.2 g/L malachite,
respectively. Therefore, it is not surprising to observe that 0.2 g/L Cu2O and 1.0 g/L malachite
also had faster HAA formation rates than 1.0 g/L CuO and 0.2 g/L malachite, respectively.
The current observation of increased HAA formation rates supports the theory of Cu-
catalyzed formation of HAA by increasing the reactivity of NOM rather than Cu-induced
chlorine decay (which would result in decreased HAA formation). Since free chlorine
degradation and HAA formation in the presence of these copper corrosion products were
reactions mediated by both aqueous and surface-bound Cu(II), further investigation is needed to
0
40
80
120
160
200
No Cu(II)and solids
CuO 1g/L Cu2O0.2g/L
Malachite0.2 g/L
Malachite1.0 g/L
HA
A9 (
µg
/L)
115
evaluate the reactivity of these solids with free chlorine and their potential to enhance HAA
formation.
Figure 5-8 HAA9 formation rates for MWTP water containing dissolved Cu(II) without solids
and solid copper corrosion products
HAA Speciation 5.3.3
Among the nine possible HAA species in MWTP water, the sum of the concentrations of
MCAA, DCAA and TCAA generally contributed to 65-93% of HAA9 independent of copper
concentration for the reaction periods examined. In addition, the concentrations of BCAA and
BDCAA were significantly higher than for other brominated species, especially at low pH.
Therefore, Figure 5-9 shows a comparison of these five more predominant HAA species
(MCAA, DCAA, TCAA, BCAA and BDCAA) after a reaction time of 100 hours under the
influence of pH in the absence and presence of 1.0 mg/L Cu(II). When Cu(II) was absent, TCAA
was the dominant species at low pH but DCAA was predominant at neutral and high pH. When
1.0 mg/L Cu(II) was present, DCAA formation consistently exceeded MCAA and TCAA
formation at all three pH levels. The MCAA concentration tended to increase with increasing pH
both with and without a spike of 1.0 mg/L Cu(II). BCAA concentrations were the same at all
three pH levels in the absence of Cu(II), but exhibited a slight increase with increasing pH in the
Cu(II)-spiked water. Conversely, BDCAA had a higher concentration at pH 6.6 than at pH 8.6 in
the absence of Cu(II), but significantly decreased and exhibited no significant difference in its
concentrations at all three pH levels in the presence of Cu(II).
10
20
30
40
50
0.0 0.2 0.4 0.6 0.8 1.0HA
A f
orm
ation r
ate
[µ
g·L
-1·ln(h
)-1]
Measured dissolved Cu(II) (mg/L)
Cu(II) without solids 0.2 g/L Cu2O1.0 g/L CuO 0.2 g/L malachite1.0 g/L malachite
116
Figure 5-9 Effect of pH on HAA speciation in the absence and presence of 1 mg/L Cu(II) at
reaction time of 100 hours for MWTP water (error bars represent standard deviation of triplicate
tests)
These observations indicate that chlorination of HAA organic precursors was influenced
by both pH and the presence of Cu(II). The observed lack of significant differences in DCAA
formation yields for the three pH values and the prevalence of TCAA at low pH in the absence of
Cu(II) are in agreement with previous studies (Reckhow and Singer, 1985; Reckhow et al., 1990;
Liang and Singer, 2003).
However, enhanced formation of DCAA with increasing pH in the presence of Cu(II)
suggests that Cu(II) would preferentially complex with DCAA precursors and increase their
reactivity with chlorine. In the presence of Cu(II), although elevated reactivity of NOM by Cu(II)
complexation (especially at high pH) led to more chlorine being involved in reactions with
NOM, the increased reactivity of DCAA precursors by Cu(II) dominated the formation of
DCAA. Furthermore, the shift of chlorine species at high pH from the potent halogenating HOCl
to less reactive OCl- should have resulted in a lesser formation of halogenated species (such as
TCAA). The different behavior of DCAA and TCAA formation under different pH conditions
suggests that their formation mechanisms or precursors may be different, especially in the
investigated water source. However, the similar pattern of pH effects for DCAA and BCAA
formation as well as for TCAA and BDCAA formation indicate that each species belonging to
the same subclass of HAA9 (i.e. dihaloacetic acids and trihaloacetic acids) had similar precursors
and reaction pathways with chlorine. All of these findings are also in agreement with other
studies (Reckhow and Singer, 1985; Liang and Singer, 2003). In terms of MCAA, Reckhow et
al. (1990) and Liang and Singer (2003) stated that pH did not impact MCAA formation.
0
20
40
60
80
100
120
pH 6.6 pH 7.6 pH 8.6 pH 6.6 pH 7.6 pH 8.6
HA
A (
µg
/L)
MCAA DCAA TCAA BCAA BDCAA
0 mg/L Cu(II) 1.0 mg/L Cu(II)
117
However, the increased MCAA formation rate at high pH that was observed in the absence and
presence of Cu(II) in this study may be due to different compositions of NOM in the water tested
in their research and in this study.
In the water samples spiked with copper corrosion solids, DCAA formation was still
consistently predominant over TCAA over the reaction period (figures are not shown). Again,
this suggests that DCAA precursors, rather than MCAA and TCAA precursors, preferentially
interacted with the dissolved copper species and/or surface-associated Cu(II) from these copper
solid corrosion products, thereby suppressing MCAA and TCAA formation. The shift of chlorine
species from HOCl to OCl- at high pH of 8.3 also made the formation of highly chlorinated HAA
species less likely (Cowman and Singer, 1996).
5.4 Summary
This is the first study that investigated catalytic potential of copper corrosion products,
including Cu(II), Cu2O, CuO, and Cu2(OH)2CO3, on chlorine decay and HAA formation. The
impacts of pH and the concentrations of these corrosion products on copper catalysis were
evaluated under controlled experimental conditions at bench scale. The following conclusions
can be drawn based on the experimental results:
a. For the three pH levels investigated, accelerated chlorine decay and HAA formation were
observed at pH 8.6 upon the addition of 1.0 mg/L Cu(II). Further investigation of
chlorine decay pathways in the presence of Cu(II) with synthetic water indicated that the
presence of dissolved and/or solid surface-associated Cu(II) would increase the reactivity
of NOM with chlorine. As a result, chlorine decay was accelerated, likely by reacting
with active Cu(II)-NOM complexes, and HAA formation was enhanced.
b. Free chlorine decayed faster and higher concentrations of HAA formed as the Cu(II)
concentration increased from 0 mg/L to 1.0 mg/L, but an absence of further changes with
increases of Cu(II) from 1.0 mg/L to 2.0 mg/L suggests that the capacity of NOM-Cu
interactions had been reached by that point.
c. Solid copper corrosion products were also observed to catalyze both chlorine decay and
HAA formation.
d. The presence of Cu(II) and its solid corrosion products led to DCAA formation
118
consistently predominating over other HAA species.
Results of this study may provide some implications for distributed water quality in
domestic plumbing systems where copper pipes are primarily installed. These pipe surfaces are
often covered by corrosion solids, and in the aqueous phase dissolved copper ions can also be
sorbed onto corrosion solids. From the present study, interactions of free chlorine and HAA
precursors with copper corrosion products will likely affect the stability of secondary
disinfectants and the fate of HAAs, especially in the premise plumbing of distribution systems.
Understanding the catalytic potential of copper on chlorine degradation and HAA formation will
be of benefit to water utilities and households for the management of their distribution systems
and water quality. Since copper concentration, pH and reaction time are the main factors to
impact the nature and extent of copper catalysis in the present study, considerations for corrosion
preventive strategies are suggested to include pH adjustment, addition of corrosion inhibitors and
periodic flushing to decrease contact time of water with pipe corrosion products to alleviate
HAA formation by copper catalysis.
5.5 References
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Alvarez-Puebla, R. A., Valenzuela-Calahorro, C., and Garrido, J. J. (2004b) Retention of Co(II),
Ni(II), and Cu(II) on a purified brown humic acid. Modeling and characterization of the
sorption process. Langmuir, 20(9), 3657-3664.
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electrochemical wet oxidation of phenol using new copper(II) tetraazamacrocycle
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Blatchley, E. R., Margetas, D., and Duggirala, R. (2003) Copper catalysis in chloroform
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119
Chun, C. L., Hozalski, R. M., and Arnold, T. A. (2005) Degradation of drinking water
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Cowman, G. A., and Singer, P. C. (1996) Effect of bromide ion on haloacetic acid speciation
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Fu, J., Qu, J. H., Liu, R. P., Qiang, Z. M., Liu, H. J., and Zhao, X. (2009b) Cu(II)-catalyzed
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122
Factors Affecting Copper Catalysis of NDMA Formation 6
from DMA in Simulated Premise Plumbing
The results of this chapter have been submitted for publication as follows:
Zhang, H., Andrews, S.A. Factors Affecting Copper Catalysis of NDMA Formation from
DMA in Simulated Premise Plumbing. Water Research.
Results from this chapter focus on the research gap “What are the impacts of metal
corrosion products on the fate of NDMA in distribution systems?”
123
Abstract
This study investigated the effects of copper, a metal commonly employed in household
plumbing systems, on N-nitrosodimethylamine (NDMA) formation from a known NDMA
precursor, dimethylamine (DMA). Copper-catalyzed NDMA formation increased with increasing
copper concentration, DMA concentration, alkalinity and hardness, but decreased with
increasing natural organic matter (NOM) concentration. pH influenced the speciation of
chloramine and the interactions of copper with DMA. The transformation of monochloramine
(NH2Cl) to dichloramine (NHCl2) and complexation of copper with DMA were proposed to be
involved in elevating the formation of NDMA by copper at pH 7.0. The inhibiting effect of
NOM on copper catalysis could be attributed to the rapid consumption of NH2Cl by NOM and/or
the competitive complexation of NOM with copper to limit the formation of DMA-copper
complexes. Hardness ions, as represented by Ca2+
, also competed with copper for binding sites
on NOM, thereby weakening the inhibitory effect of NOM on NDMA formation. Common
copper corrosion products also participated in these reactions but in different ways. Aqueous
copper released from malachite [Cu2CO3(OH)2] was shown to promote NDMA formation while
NDMA formation decreased in the presence of CuO, possibly due to the adsorption of DMA.
Keywords: Copper; NDMA; NOM; premise plumbing
124
6.1 Introduction
With the implementation of the Stage 2 Disinfectant/Disinfection Byproduct Rule
D/DBPR, combined chlorine has received increasing attention as a secondary disinfectant in
distribution systems. The formation of monochloramine (NH2Cl) generally dominates over other
species of combined chlorine at a pH value of approximately 8 and at a Cl2/N molar ratio of less
than 1:1. Dichloramine (NHCl2) is formed at higher Cl2/N ratios (molar ratio >1:1) or at low pH
values (pH <5), and its formation is strongly catalyzed by the presence of acetic acid, phosphate,
carbonate, and silicate species (Schreiber and Mitch, 2007; Valentine et al., 1988). It is well
established that fewer halogenated disinfection byproducts (DBPs) are formed upon
chloramination than chlorination. However, the application of chloramines also causes some
unintended changes in water quality, including increased formation of nitrosamines (Choi et al.,
2002; Najm and Trussell, 2001) and elevated metal corrosion rates (Boyd et al., 2008).
One of the typical reaction pathways for NDMA formation is from the reaction of NH2Cl
and dimethylamine (DMA). Two steps are included in this mechanism: a nucleophilic
substitution of DMA by NH2Cl to form unsymmetrical dimethylhydrazine (UDMH) and the
subsequent oxidation of this UDMH intermediate to NDMA (Choi and Valentine, 2002a; Mitch
et al., 2003). The rate of UDMH formation is slow and increases with an increase in pH (Kim
and Clevenger, 2007), whereas the UDMH oxidation occurs nearly instantaneously but with low
yields (<1%) (Mitch and Sedlak, 2002). In addition, Schreiber and Mitch (2006, 2007) proposed
two other NDMA formation pathways involving NHCl2: a relatively slow reaction of NHCl2
with amine precursors in the presence of dissolved oxygen, and a fast reaction involving a
reactive breakpoint chlorination intermediate (HNO) which is produced from the hydrolysis of
NHCl2. Normally, NHCl2 can be minimized by controlling pH (> 8.5) and Cl2/N molar ratios
(<1) (Schreiber and Mitch, 2006).
Drinking water guidelines for nitrosamines have been under recent development in
Canada and the United States. Health Canada has established a guideline for NDMA in drinking
water at a Maximum Acceptable Concentration (MAC) of 40 ng/L (Health Canada, 2011). The
Ontario Ministry of the Environment, Canada, has set a standard for NDMA of 9 ng/L (MOE,
2002), and the California Department of Public Health has set notification levels of 10 ng/L for
125
each of NDMA, N-nitrosodiethylamine (NDEA) and N-nitrosodi-n-propylamine (NDPA)
(CDPH, 2009).
Copper is commonly used in pipes and copper alloys that are found in household
plumbing. Copper corrosion products in drinking water systems include copper ions [Cu(II)],
cuprite [Cu2O], tenorite [CuO], cupric hydroxide [Cu(OH)2], and malachite [Cu2CO3(OH)2]
(Xiao et al., 2007). The catalytic potential of various copper species in many reaction processes
has been widely investigated (Onuchukwu, 1994; Pintar et al., 1997; Paidar et al., 1999). Some
studies have demonstrated copper catalysis on free chlorine and NH2Cl degradation as well as
THM and HAA formation (Blatchley et al., 2003; Fu et al., 2009b; Fu et al., 2009c; Li et al.,
2007). Although Schreiber and Mitch (2006) reported elevated NDMA formation in the presence
of copper, details concerning factors that impact copper catalysis during NDMA formation are
not available.
The most likely mechanism of copper catalysis during THM and HAA formation has
been speculated to involve interactions of Cu(II) with natural organic matter (NOM) and a
subsequent increased reactivity of NOM with chlorine. NOM is composed of an extremely
diverse group of materials that include carboxyl groups, some phenolic groups, alcohol groups,
methoxyl groups, ketones, and aldehydes (Reckhow et al., 1990). The binding affinity of NOM
with Cu(II) has been reported to depend on the types of functional groups involved, pH, copper
concentration, and the concentrations of Ca2+
and Mg2+
(Brown et al., 1999; Lu and Allen, 2002;
Rate et al., 1993). For example, according to Lu and Allen (2002), hardness ions (Ca and Mg)
had competitive impacts on Cu-NOM complexation, although the binding affinities of Ca and
Mg are much weaker than that of Cu. They also reported that competitive impacts of hardness
ions on Cu complexation increased with increasing pH and hardness concentration, but these
impacts became less significant at successively higher hardness concentrations.
The objective of this study was to investigate the catalytic potential of Cu(II) on NDMA
formation under controlled experimental conditions by performing simulated distribution system
(SDS) tests. Investigated factors that were theorized to affect copper’s impact on NDMA
formation from DMA included copper concentrations, DMA concentrations, alkalinity, NOM
and hardness. Since copper is widely used in domestic plumbing systems, understanding the
impacts of copper and its corrosion products on the formation of NDMA may help utilities and
126
households to minimize the potential for NDMA contamination in their systems by suggesting
preventive strategies.
6.2 Materials and methods
Chemicals and Materials 6.2.1
All chemicals used in this study were ACS grade or higher. High quality Milli-Q water
produced from a Millipore Milli-Q UV plus ultrapure water system (Millipore, Mississauga, ON)
was used for all water blanks and chemical reagent solutions. Dissolved Cu(II) was added in the
form of CuSO4 (1000 mg/L stock solution). Phosphate buffer (H2PO4-/HPO4
2-) and borate buffer
(H3BO3/NaOH) solutions were prepared and added to test solutions at 1 mM concentration to
control reaction solution pH at desired levels. The impacts of different buffers at a concentration
of 1 mM on NDMA formation were negligible according to Schreiber and Mitch (2006). Sulfuric
acid (H2SO4, 50%) and sodium hydroxide (NaOH, 20%) were also used for pH adjustment. A
sodium carbonate (Na2CO3, 300 mM) solution was used to adjust alkalinity. A calcium chloride
(CaCl2, 1000 mM) solution was prepared for hardness tests. Suwannee River natural organic
matter (SR-NOM), purchased from the International Humic Substances Society, was added to
Milli-Q water to adjust NOM concentration. Dimethylamine (DMA, 2.0 M in methanol) was
purchased from Sigma-Aldrich. In all of the tests except for those evaluating effects of DMA
concentrations, DMA at a concentration of 11.2 µg/L (250 nM) was dosed to ensure the
formation of NDMA at measureable levels. L-Ascorbic acid (≥ 99.0%, Sigma Aldrich) at 200
mg/L was applied to quench total chlorine residuals before NDMA extraction.
The free chlorine dosing solution (approximately 3500 mg/L as Cl2) was made by
diluting a concentrated solution of sodium hypochlorite (NaOCl, 6%, VWR) in Milli-Q water.
The resulting solution was standardized by further dilution with Milli-Q water and then
measurement using the DPD colorimetric method with a Hach DR 2700 spectrophotometer. The
total chlorine dosing solution, containing at least 90% monochloramine (NH2Cl), was preformed
by adding the free chlorine dosing solution to a 2.7 g/L (50 mM) ammonium chloride solution in
a well-stirred 250 mL amber bottle at a Cl2/N molar ratio of 0.8:1.
127
Experimental Procedures 6.2.2
All of the experiments were performed in 1 L amber bottles with PTFE-lined caps and at
room temperature (20±1 ºC). The reaction bottles were made chlorine demand free before use.
In the tests to evaluate effects of copper concentration on NDMA formation, copper
sulphate was spiked to provide copper ion concentrations ranging from 0 to 1.0 mg/L to be
compatible with Drinking water guidelines for copper – maximum acceptable concentration 1.0
mg/L (Health Canada, 2009). Malachite [Cu2CO3(OH)2] and tenorite (CuO) were selected as the
solid forms of copper corrosion products of interest according to the results of MINEQL+
version 4.5 (Environmental Research Software, Hallowell, Maine, USA) simulations. They were
added into test solutions in the form of powders and mixed with test water thoroughly using a stir
bar. After 24 hours, water samples dosed with corrosion products were filtered through 0.2 µm
Nylaflo® Nylon membrane filter paper (Pall Corporation) to separate the solid copper corrosion
products from the aqueous solution. Dissolved copper in the filtrate was measured by flame
atomic absorption spectrometry (FAAS). To account for any influence from filter paper and
filtration performance on NDMA concentrations, control samples without the addition of solid
copper corrosion products were also filtered under similar conditions.
To evaluate the impacts of NOM on copper solubility and to determine if DMA-Cu
complexes played a role in copper catalysis, copper sulphate was spiked in the presence of
different concentrations of NOM and/or DMA. The resulting solutions were equilibrated for 6
hours and then filtered through 0.2 µm Nylaflo® Nylon membrane filter paper, and copper
concentrations in the filtrate were measured by FAAS.
All experiments were performed at pH 7.0±0.1 (buffered by 1 mM H2PO4-/HPO4
2-
solution) in Milli-Q water except when evaluating pH effects, and the copper level was
maintained at 1 mg/L except in the tests where effects of copper concentrations were evaluated.
An initial NH2Cl concentration of 2.3±0.1 mg/L was applied for all of the tests, chosen to be
typical of that which is observed in drinking water entering distribution systems.
The concentration and speciation of chloramine in the absence and presence of copper
were determined according to Schreiber and Mitch (2005, 2006). A CE 3055 Reflectance
Spectrophotometer (Cecil Instruments, England) was used. To ensure the measurability of
128
NH2Cl and dichloramine (NHCl2) via spectrophotometric analysis, a preformed NH2Cl stock
solution at a concentration of 130 mg/L was tested. The solution pH was controlled at pH 6.7 and
7.0, respectively. After 20 hours reaction, NH2Cl and NHCl2 were distinguished and quantified
by monitoring absorbance at their respective λmax (λmax, NH2Cl= 2 , λmax, NHCl2= 295) and solving
simultaneous equations.
NDMA extract preparation and measurement were carried out according to Taguchi et al.
(1994). Isotope-labeled surrogate (d6-NDMA) was used to correct for interferences from
nitrosamine extraction processes and matrix effects. The instrument used for nitrosamine
analysis was a Varian 4000 GC/MS operated in chemical ionization mode.
6.3 Results and discussion
Since alkalinity, NOM, hardness and pH may affect copper speciation and solubility
(Broo et al., 1998; Korshin et al., 1996), the effects of these water quality parameters on NDMA
formation from DMA were all examined in this study. The impacts of DMA concentrations on
NDMA formation and copper solubility as well as its major complexes/precipitates were also
tested.
Effect of Copper Concentrations 6.3.1
Figure 6-1 illustrates the observed NDMA formation when varying the copper sulphate
concentrations (expressed as mg/L total Cu). NDMA formation increased approximately linearly
(31 to 104 ng/L) when spiked Cu(II) concentrations were increased from 0 mg/L to 1.0 mg/L,
suggesting some participation or possible catalytic effect of copper during NDMA formation
from DMA.
Separate NDMA formation kinetics tests also showed that less than half the amount of
NDMA would form in the absence of copper than in the presence of 1.0 mg/L copper (Figure
6-2), providing further support to the theory that copper may enhance or catalyze NDMA
formation.
129
Figure 6-1 NDMA formation with increasing copper concentrations from added CuSO4; pH 7.0,
initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA spiked, 24 hours, error bars indicate the measured
maximum and minimum values (n=2)
Figure 6-2 NDMA formation kinetics in the absence and presence of copper; pH 7, Cu(II) 1
mg/L spiked, initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA, error bars indicate the measured
maximum and minimum values (n=2)
Effects of DMA Concentration and Cu-NH2Cl Interactions 6.3.2
Figure 6-3 shows a comparison of the impacts of DMA concentrations (varying from 0 to
1800 µg/L) on NDMA formation in the absence and presence of 1.0 mg/L spiked Cu(II).
Although this range of concentrations extends well above the range of any NDMA precursors
that would be expected in typical drinking water sources, the test shows the extent of the effects
of copper presence, even at a relatively low concentration of 1.0 mg/L Cu(II). In the absence of
copper, NDMA formation increased with increasing DMA concentrations, and the yield of
0
30
60
90
120
0 5 10 15 20 25
ND
MA
(ng/L
)
Time (hours)
1 mg/L Cu(II)
0 mg/L Cu(II)
0
30
60
90
120
0.0 0.1 0.2 0.3 0.4 0.5 0.7 1.0
ND
MA
(ng/L
)
Copper from CuSO4 (mg/L)
130
NDMA at 24 hours was 2670 ng/L at a DMA concentration of 1800 µg/L (approximately 0.15
±0.01% yield, which is typical for this precursor). In the presence of copper, NDMA formation
was enhanced by 2 to 4 times compared with that in the absence of copper. These data further
confirmed the catalytic impacts of copper on NDMA formation that were shown in Figures 6-1
and 6-2. The inset plot in Figure 6-3 shows NDMA yields at DMA concentrations ranging from
0 to 11.2 µg/L. Comparable NDMA yields were obtained for the samples that had similar
conditions to those of Figure 6-1 (e.g. 11.2 µg/L DMA and 1 mg/L copper produced 107 ng/L vs.
104 ng/L NDMA), providing further evidence of the excellent reproducibility and reliability of
the results in this study.
Figure 6-3 NDMA formation with DMA concentrations; pH 7.0, initial NH2Cl 2.3±0.1 mg/L,
Cu(II) 1 mg/L spiked, 24 hours, d error bars indicate the measured maximum and minimum
values (n=2)
In addition, as shown in Figure 6-3, the yields of NDMA in the absence of copper were
approaching a maximum of approximately 2600 ng/L at the molar ratio of DMA (1800 µg/L, 40
µM) to NH2Cl (2.4 mg/L, 34 µM) of approximately 1.2:1, which is close to the 1:1 molar ratio
for the maximum NDMA formation observed by Choi and Valentine (2002b). As they proposed,
a further increase in the DMA:NH2Cl ratio would make the amount of NH2Cl available to
oxidize the UDMH intermediate to be rapidly depleted due to chlorine transfer. In the current
research and in the presence of copper, a maximum NDMA yield occurred at a DMA
concentration of approximately 900 µg/L (20 µM) with the DMA:NH2Cl ratio of only 0.6. This
shift of the maximum DMA:NH2Cl ratio in the presence of copper is proposed to be associated
with copper’s interactions with NH2Cl, given that copper has been shown to enhance the
decomposition of NH2Cl to NHCl2 under acidic conditions (Fu et al., 2009a; Fu et al., 2009c)
0
2000
4000
6000
8000
0 400 800 1200 1600 2000
ND
MA
(n
g/L
)
DMA (µg/L)
0 mg/L Cu(II)
1 mg/L Cu(II)
0
50
100
150
0 5 10 15
107 ng/L
131
and NHCl2 has been previously reported to increase NDMA formation (Schreiber and Mitch,
2006 and 2007).
Separate tests utilizing the spectrophotometric method described by Schreiber and Mitch
(2005, 2006) were performed to confirm the proposed shift in the relative amounts of NH2Cl and
NHCl2 by determining the concentration and speciation of chloramine in the absence and
presence of copper. As shown in Figure 6-4, the addition of 1 mg/L copper at pH 7.0
significantly increased NHCl2 formation (by 2 fold) while the NH2Cl concentration was
decreased by 45%. The effect was even further enhanced at a slightly lower pH (pH 6.7), and
will be discussed again in Section 6.3.7. Thus, the catalyzed formation of NHCl2 by copper at pH
7.0 could be one reason for the enhanced NDMA formation observed in Figure 6-3. Furthermore,
the decomposition of NH2Cl by copper limited the amount of NH2Cl which could participate in
the reactions for oxidizing the UDMH intermediate to NDMA, thereby shifting the concentration
of DMA for the maximum formation of NDMA from 1800 µg/L in the absence of copper to 900
µg/L in the presence of copper (Figure 6-3).
Figure 6-4 Chloramine speciation in the absence and presence of copper at pH 6.7 and 7, Milli-Q
water
Cu-DMA Complexation 6.3.3
In addition to the influence of increased NHCl2 on enhancing the formation of NDMA,
the formation of DMA-Cu complexes was also hypothesized to be able to increase the
conversion of DMA to NDMA. The ability of Cu(II) to complex with organic matter has been
0
30
60
90
120
0
20
40
60
80
100
120
6.7 7.0 6.7 7.0
Su
m o
f N
H2C
l a
nd
N
HC
l 2(m
g/L
as C
l 2)
Chlo
rin
e (
mg
/L a
s C
l 2)
pH
NH₂Cl NHCl₂ Sum of NHCl₂ and NH₂Cl
Without copper With copper
132
reported to be essential for copper to act as a catalyst (Bansal et al., 2008). Aside from
carboxylic and phenolic groups, copper has the potential to complex with the lone pair of
electrons on the central nitrogen of DMA (Croue et al., 2003; Tuschall and Brezonik, 1980;
Westerhoff and Mash, 2002). The complexation of copper with organic nitrogen is moderately
stable, with conditional stability constants ranging from 1.6 106 to 1.3 10
7 (Ko and Lee, 2010;
Tuschall and Brezonik, 1980). Therefore, to detect the occurrence of DMA-Cu complexation,
five concentrations of DMA (0, 45, 360, 702 and 3510 µg/L) were dosed in Milli-Q water
containing 1.0 mg/L copper at pH 7.0, and the amount of copper remaining in dissolved form
was measured. If the complexation between the nitrogen of DMA and copper occurred, the
solubility of copper was expected to increase. As shown in Figure 6-5, the filtered copper
concentration was only 0.3 mg/L in the absence of DMA (0 µg/L), and increased to 0.4 and 0.8
mg/L Cu at DMA concentrations of 45 and 360 µg/L, respectively. Further additions of DMA to
702 and 3510 µg/L did not change the filtered copper concentration significantly. Regardless, the
increased solubility of copper for solutions with DMA at concentrations below 360 µg/L strongly
suggests the complexation of the nitrogen in DMA with copper. This complexation can
essentially increase the reactivity of DMA, thereby enhancing NDMA formation.
Figure 6-5 Dissolved copper concentrations with varying DMA; filtered by 0.2 µm Nylaflo®
Nylon membrane filter paper, pH 7, Cu(II) 1 mg/L spiked, 6 hours, error bars indicate the
measured maximum and minimum values (n=2)
Effect of Alkalinity 6.3.4
The effect of alkalinity was examined in this study because copper can form carbonate-
containing species which may play a role in copper’s participation in NDMA-forming reactions.
0.0
0.2
0.4
0.6
0.8
1.0
0 45 360 702 3,510
Dis
solv
ed c
opper
(mg/L
)
DMA (µg/L)
133
Figure 6-6 illustrates the yields of NDMA at different alkalinity levels along with the likely
major copper species determined by MINEQL+ version 4.5. For alkalinity values below 50 mg/L
as CaCO3, NDMA concentrations remained approximately constant at 94~100 ng/L. For
alkalinity values greater than 50 mg/L as CaCO3, NDMA formation increased approximately
linearly from 98 to 225 ng/L (R2
=0.9), suggesting a possible promoting effect of alkalinity on
copper catalyzed NDMA formation.
Figure 6-6 NDMA formation with increasing alkalinity and copper speciation as a function of
alkalinity (determined by MINEQL+ version 4.5); pH 7.0, initial NH2Cl 2.3±0.1 mg/L, Cu(II) 1
mg/L spiked, 11.2 µg/L DMA, 24 hours, duplicate
This apparent promotional effect of alkalinity may have been more related to the copper
species present than to the alkalinity itself. According to MINEQL+ simulations, at pH 7.0 the
spiked copper was present mainly in the form of precipitates (> 85%). For alkalinity values
below 50 mg/L as CaCO3, which generally would be much lower than observed in distribution
systems, Cu3(PO4)2 was the primary form of solids due to the addition of phosphate buffer; for
alkalinity values above 50 mg/L as CaCO3, malachite [Cu2(OH)2CO3] was the dominant solid.
The rest of the copper was present in the dissolved form as free Cu(II) and soluble complexes,
and the concentration of these dissolved components increased with increasing alkalinity.
When Cu3(PO4)2 was present at concentrations of between 0.92 and 0.97 mg/L as copper,
again, only at very low alkalinities and so not likely to be seen in distribution systems, NDMA
yields remained constant, indicating that such low alkalinity values and/or Cu3(PO4)2 had no
significant impacts on NDMA formation from DMA. In the presence of malachite, given that
0
50
100
150
200
250
0.0
0.2
0.4
0.6
0.8
1.0
0 20 35 50 100 150 200 300
ND
MA
(n
g/L
)
Cop
pe
r sp
ecia
tio
n (
mg
/L)
Alkalinity (mg/L as CaCO3)
Aqueous copper PrecipitationNDMA
Cu3(PO
4)2 Cu
2(OH)
2CO
3
134
NDMA formation increased with the increase in the corresponding aqueous copper
concentrations, it follows that the copper catalysis of NDMA formation was essentially a
reaction mediated by aqueous copper released from the malachite. Separate experiments were
performed to test NDMA formation and copper release in the presence of pre-dosed malachite.
The yields of NDMA were increased from 24 ng/L in the absence of malachite to 144 and 171
ng/L after dosing 0.1 and 1.0 g/L malachite, respectively, whereas the released copper
concentrations in the presence of 0.1 and 1.0 g/L malachite were 0.2 and 0.6 mg/L, respectively.
These results support the theory that aqueous copper released from malachite could catalyze
NDMA formation.
Effect of NOM 6.3.5
As discussed in Section 6.3.3, copper can complex with DMA and thus increase DMA
reactivity to form NDMA (Figure 6-5). In all natural waters, natural organic matter (NOM) is
ubiquitous and it also has a strong affinity for metal ions (Stevenson, 1994). Therefore, to
demonstrate the competitive impacts of NOM on the interactions between DMA and copper and
the subsequent NDMA formation, a range of concentrations of Suwannee River NOM (SR-
NOM) were dosed into DMA-spiked water in the presence of 1 mg/L copper. NDMA formation
results are shown in Figure 6-7, along with corresponding NH2Cl residual and dissolved copper
concentrations after 24-h reaction time. In the absence of SR-NOM, 111 ng/L NDMA formed
which is compatible to the yields obtained in previous experiments (104 and 107 ng/L). NDMA
Figure 6-7 NDMA, NH2Cl residual and dissolved copper with increasing SR-NOM
concentrations; pH 7.0, initial NH2Cl 2.3±0.1 mg/L, Cu(II) 1 mg/L spiked, 11.2 µg/L DMA, 24
hours, error bars indicate the measured maximum and minimum values (n=2)
0
30
60
90
120
0.0 1.0 2.0 3.0 4.0 5.0
ND
MA
(n
g/L
)
SR-NOM (as mg/L TOC)
0.0
0.2
0.4
0.6
0.8
1.0
0.0
0.5
1.0
1.5
2.0
2.5
0.0 1.0 2.0 3.0 4.0 5.0
SR-NOM (mg/L as TOC)
Dis
so
lve
d C
u(I
I) (
mg
/L)
NH
2C
l re
sid
ua
l (m
g/L
)
NH₂Cl residual
Dissolved Cu(II)
135
concentrations decreased sharply for SR-NOM concentrations up to 1.0 mg/L as TOC, and then
more gradually decreased to 26 ng/L with further increases in SR-NOM concentrations to 4.1
mg/L as TOC. Figure 6-7 also shows that increases in SR-NOM concentrations corresponded to
NH2Cl residual decreases, whereas the dissolved copper concentrations increased.
There are two possible explanations for the observed decreases in NDMA formation.
First, SR-NOM likely formed strong complexes with copper, as indicated by the corresponding
increased dissolved copper concentrations with increasing SR-NOM concentrations. This
complexation, therefore, would likely compete with the interactions of copper with DMA and
subsequently decrease the catalytic effects of copper on NDMA formation. In addition, SR-NOM
also competes with DMA for NH2Cl, as indicated by the decreased NH2Cl residual concentration
upon increased SR-NOM concentrations, and NH2Cl is a rate-limiting factor for NDMA
formation (Chen and Valentine, 2006). For both of these reasons, it was reasonable to have
observed decreased NDMA formation after the addition of SR-NOM.
Effect of Hardness 6.3.6
The impacts of hardness ions on copper catalysis during NDMA formation have also
been investigated in this study given that these ions may compete with copper for the available
binding sites on NOM or DMA. Due to the abundance of hardness ions in most natural waters,
the impacts of these ions on copper catalysis may be an important aspect to consider when
determining NOM (and/or DMA)-Cu complexation. Since calcium is the dominant hardness-
related ion in most natural waters, calcium (Ca2+
) was chosen to be representative of the hardness
ions in this study.
Figure 6-8 displays the NDMA yields and corresponding dissolved copper concentrations
obtained at hardness concentrations varying from 0 to 200 mg/L as CaCO3 in the presence of 4.1
mg/L SR-NOM as TOC. NDMA formation increased with increasing hardness while the
dissolved copper concentrations exhibited the opposite trend. This suggests that Ca2+
can
compete with copper for the available binding sites on SR-NOM, thereby decreasing the amount
of copper complexed with NOM (NOM-Cu complexation was discussed in Section 6.3.5). It
follows that copper-catalyzed NDMA formation would be increased in hard waters due to the
competitive effects of Ca2+
on NOM-Cu complexation, freeing up copper ions to participate in
NDMA formation reactions.
136
Figure 6-8 Variation of NDMA formation and dissolved copper concentrations with increasing
hardness; pH 7.0, Cu(II) 1 mg/L spiked, initial NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA spiked,
SR-NOM 4.1 mg/L as TOC, 24 hours, error bars indicate the measured maximum and minimum
values (n=2)
Effect of pH and the Role of NHCl2 in NDMA Formation 6.3.7
This study also investigated the impact of pH on copper catalysis during NDMA
formation since pH can affect the speciation of chloramine and copper as well as the interactions
of copper with DMA and organic matter. Figure 6-9 summarizes the observed NDMA yields and
corresponding pseudo-first-order NH2Cl decay constants for tests performed in the absence and
presence of 1.0 mg/L Cu(II) at pH levels that span a range found in distribution systems. In the
absence of added Cu(II), NDMA formation from DMA increased with increasing pH, which is in
agreement with previous findings (Kim and Clevenger, 2007). When 1 mg/L Cu(II) was present,
NDMA formation was enhanced at pH 6.7, 7.0 and 8.0 by factors of 8.0, 2.1 and 2.3,
respectively. At pH 9.0, however, NDMA formation was only approximately 80% of that formed
in the absence of Cu(II).
The observed increased NDMA formation under slightly acidic or neutral conditions (pH
6.7 and 7.0), as shown in Figure 6-9, has been proposed to be due to a greater formation of
NHCl2, as discussed in Section 6.3.2. As estimated by NH2Cl pseudo-first-order decay constants
at four pH levels (shown in Figure 6-9), regardless of the presence of Cu(II), NH2Cl significantly
decayed at both pH 6.7 and pH 7.0, but remained approximately constant over the test period of
48 hours at pH 8.0 and pH 9.0. It is known that NH2Cl is more stable at pH 8.0 and that it
0.0
0.2
0.4
0.6
0.8
1.0
0
10
20
30
40
50
0 50 100 150 200
Dis
so
lve
d C
u(I
I) (
mg
/L)
ND
MA
(n
g/L
)
Hardness (mg/L as CaCO3)
NDMA Dissolved Cu(II)
137
decomposes to NHCl2 under even slightly acidic conditions (Valentine et al., 1988). In this
study, separate tests using the spectrophotometric method described in Section 6.2.2 confirmed
an enhancement in the formation of NHCl2 by copper under acidic and neutral conditions (Figure
6-4). Therefore, the catalyzed formation of NHCl2 by Cu(II) could be one reason for the
observed enhanced NDMA formation in the presence of Cu(II), especially at pH 6.7.
Figure 6-9 NDMA formation at four pH levels in the absence and presence of 1 mg/L Cu(II) in
Milli-Q water; 24 hour, NH2Cl 2.3±0.1 mg/L, 11.2 µg/L DMA, error bars indicate the measured
maximum and minimum values (n=2)
Since an enhanced conversion of NH2Cl to NHCl2 by Cu(II) does not occur in basic
solutions, as evidenced in this study and by Fu et al. (2009a, 2009c), the increased NDMA
formation observed at pH 8.0 in the presence of Cu(II) (Figure 6-9) was likely attributed to other
factors. For example, the complexation of copper with the central nitrogen of DMA was
discussed in Section 6.3.3, and Ko and Lee (2010) have reported the increased potential for
nitrogen to chelate Cu(II) under alkaline conditions relative to acidic conditions. Therefore,
increasing the pH from 7.0 to 8.0 was expected to form more DMA-Cu complexes, thereby
increasing the reactivity of DMA with NH2Cl and subsequent NDMA formation. The observed
NDMA formation results support this theory.
However, the results of separate filtration tests (similar to those described in Section
6.3.3) with five concentrations of DMA at pH 8.0 showed that total dissolved copper
concentrations were only increased from 0.02 mg/L in the absence of DMA to 0.05 mg/L at a
DMA concentration of 702 µg/L. This is in contrast to the results at pH 7.0 for which at least 0.8
0.000
0.003
0.006
0.009
0.012
0
100
200
300
400
500
600
6.7 7.0 8.0 9.0
Pseudo 1
st ord
er
decay c
onsta
nt (h
-1)
ND
MA
(n
g/L
)
pH levels
NDMA 0 mg/L Cu(II) NDMA 1 mg/L Cu(II)
NH₂Cl 0 mg/L Cu(II) NH₂Cl 1 mg/L Cu(II)
604 ng/L
138
mg/L copper was dissolved at the same concentration of DMA (Figure 6-5). This insignificant
formation of DMA-Cu complexes at pH 8.0 suggests that the increased formation of NDMA at
pH 8.0 is more likely a solid surface mediated reaction, and that the formation of DMA-Cu
complexes did not contribute significantly to NDMA formation at this slightly basic pH.
At pH 9.0, the yield of NDMA in the presence of Cu(II) decreased by 23% compared
with that in the absence of Cu(II). The filtration tests described in Section 6.3.3 were also
performed at pH 9.0 to identify the likely formation of DMA-Cu complexes, and the filtered
copper concentrations before and after the addition of DMA remained below Method Detection
Limit (MDL) of FAAS (0.009 mg/L). This indicates that the majority of the spiked copper was
present in the solid form (CuO, determined by MINEQL+) so DMA could not have formed
significant amount of soluble complexes with copper at pH 9.0. Since DMA has a pKa of 10.7
and NDMA is a neutral compound, it is likely that either DMA or NDMA exhibited some
adsorption to CuO at pH 9.0. In a confirmatory experiment in which NDMA was formed from
DMA in contact with a suspension of CuO at concentrations of 0.2 and 1.0 g/L, NDMA yields
decreased with increasing amounts of CuO. Further investigation is needed to describe the
adsorption behavior of DMA and/or NDMA onto CuO.
6.4 Summary
This is the first study that demonstrated the catalytic potential of copper to enhance
NDMA formation from DMA. The impacts of copper concentrations, DMA concentrations,
alkalinity, SR-NOM, hardness, and pH on copper catalysis were evaluated under controlled
experimental conditions at bench scale. The following conclusions can be drawn based on the
experimental results:
a. Complexation of copper with DMA and the subsequent increased reactivity of DMA with
NH2Cl were suggested to be essential for copper to catalyze NDMA formation.
b. NDMA formation was inhibited by the presence of SR-NOM likely because SR-NOM
may compete with DMA for NH2Cl and/or may strongly complex with copper, thereby
limiting the interactions of copper with DMA.
c. The addition of hardness ions (Ca2+
) weakened the complexation of SR-NOM with
copper by competing with copper for binding sites on SR-NOM, explaining the increased
139
formation of NDMA with increasing hardness.
d. pH had a complex impact on copper catalysis during NDMA formation, influencing the
speciation of chloramine and the interactions of copper with DMA.
e. Aqueous copper released from malachite [Cu2CO3(OH)2] was shown to promote NDMA
formation while NDMA formation decreased in the presence of CuO, possibly due to the
adsorption of DMA
This study provides information that should be considered in relation to the formation
and fate of NDMA in premise plumbing where copper pipes are primarily installed. This
improved understanding about the factors affecting NDMA formation in copper pipes will be
useful in developing strategies to control copper corrosion and reduce NDMA formation. For
example, given the demonstrated impacts of dissolved copper concentrations and retention time
on NDMA formation, the addition of phosphate-based corrosion inhibitors to control copper
corrosion are suggested. Flushing household lines before using water has been recommended to
minimize exposure to dissolved copper. The present research suggests that periodic flushing to
decrease contact time of the water with the pipe material will also help minimize exposure to
NDMA. However, further investigation is needed to evaluate the impacts of real water matrices,
including the complicated impacts of pH, alkalinity, hardness and NOM combinations with
different water compositions on the speciation and solubility of copper and the subsequent
copper catalysis during NDMA formation.
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141
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corrosion. Journal American Water Works Association, 88(7), 36-47.
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dissolved organic matter (DOM) - link to acidic moieties of DOM and competition by Ca
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dimethylamine during chlorination. Environmental Science & Technology, 36(4), 588-
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L. (2003) N-nitrosodimethylamine (NDMA) as a drinking water contaminant: A review.
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Najm, I., and Trussell, R. R. (2001) NDMA formation in water and wastewater. Journal
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Decomposition of Hydrogen-Peroxide. Materials Chemistry and Physics, 37(2), 129-131.
Paidar, M., Rousar, I., and Bouzek, K. (1999) Electrochemical removal of nitrate ions in waste
solutions after regeneration of ion exchange columns. Journal of Applied
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Pintar, A., Bercic, G., and Levec, J. (1997) Catalytic liquid phase oxidation of aqueous phenol
solutions in a trickle-bed reactor. Chemical Engineering Science, 52(21-22), 4143-4153.
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Dissociation Kinetics to Factors Influencing Complex Stability and Macromolecular
Conformation. Environmental Science & Technology, 27(7), 1408-1414.
Reckhow, D. A., Singer, P. C., and Malcolm, R. L. (1990) Chlorination of Humic Materials - by-
Product Formation and Chemical Interpretations. Environmental Science & Technology,
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formation during chloramination. Environmental Science & Technology, 39(10), 3811-
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Schreiber, I. M., and Mitch, W. A. (2006) Nitrosamine formation pathway revisited: The
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143
Schreiber, I. M., and Mitch, W. A. (2007) Enhanced nitrogenous disinfection byproduct
formation near the breakpoint: Implications for nitrification control. Environmental
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Wiley, New York.
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Determination of N-Nitrosodimethylamine by isotope-dilution, high-resolution mass-
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Valentine, R. L., Jafvert, C. T., and Leung, S. W. (1988) Evaluation of a chloramine
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144
Effects of Pipe Materials, Orthophosphate, and Flow 7
Conditions on Chloramine Decay and NDMA Formation in
Modified Pipe Loops
The results of this chapter have been submitted for publication as follows:
Zhang, H., Andrews, S.A. Effects of Pipe Materials, Orthophosphate, and Flow Conditions
on Chloramine Decay and NDMA Formation in Modified Pipe Loops. Journal of Water
Supply: Research and Technology – Aqua.
Results from this chapter focus on the research gap “How do flow conditions potentially
affect secondary disinfectant stability and NDMA formation?”
145
Abstract
Secondary disinfectants experience temporal and spatial degradation in distribution
systems, and the application of combined chlorine (mainly monochloramine) can potentially
increase the formation of nitrosamine disinfection byproducts. To date, the literature concerning
the fate of nitrosamines in full-scale water mains and household plumbing is limited and
inconsistent. This study examined the effects of pipe materials (iron, copper, lead and PVC) and
orthophosphate on chloramine decay and NDMA formation under stagnant, laminar and
turbulent flow conditions using modified pipe loops. Generally, turbulent conditions increased
the released metal concentrations compared with laminar conditions regardless of
orthophosphate addition. Orthophosphate did not statistically affect iron corrosion rates, but
effectively reduced released copper concentrations under both laminar and turbulent conditions.
Chloramine presence was a rate-limiting factor for NDMA formation, and its decay rate
generally increased with increasing flow velocity. Orthophosphate increased chloramine decay in
contact with iron by increasing nitrite formation, but decreased chloramine decay in contact with
copper and lead by reducing metal corrosion. Copper consistently catalyzed NDMA formation
from DMA under laminar flow conditions. Iron also catalyzed NDMA formation but only under
turbulent conditions. Orthophosphate increased the catalytic effects of iron by modifying the
properties of the associated suspended particles, but decreased copper catalysis by reducing the
dissolved copper concentrations.
Keywords: Chloramine; NDMA; Iron; Copper; Lead; Modified pipe loop
146
7.1 Introduction
The application of a secondary disinfectant in distribution systems is to maintain the
microbial stability of the treated water. However, secondary disinfectants may experience
temporal and spatial degradation in distribution systems due to chemical and biological reactions
that occur in the bulk water and on the pipe wall. Therefore, the overall disinfectant decay within
a pipe is usually postulated as a pseudo-first-order reaction consisting of parallel reactions
occurring in the bulk flow and on the pipe wall (Rossman et al., 1994). Generally, both bulk
water and pipe wall disinfectant decay rates increase with increasing flow velocity as a result of
an increased mass transfer rate of disinfectant residual to the pipe surface, and/or an increased
release rate of corrosion products from the pipe surface (Digiano and Zhang, 2005; Hallam et al.,
2002; Mutoti et al., 2007; Westbrook and Digiano, 2009).
Recently, combined chlorine has received growing attention by water utilities that are
challenged to comply with the Stage 2 Disinfectant/Disinfection Byproducts Rule D/DBPR. The
advantages of applying combined chlorine as an alternative to free chlorine include a minimized
production of trihalomethances (THMs) and haloacetic acids (HAAs), and a relatively stable
residual in distribution systems. However, combined chlorine (mainly monochloramine) can
potentially increase the formation of nitrosamines (Choi et al., 2002; Najm and Trussell, 2001),
of which N-nitrosodimethylamine (NDMA) is the major species of interest. The occurrence of
nitrification and elevated lead release have also been observed in chloraminated distribution
systems (Vasquez et al., 2006; Wilczak et al., 1996).
Considerable effort has been made to investigate factors affecting disinfection byproduct
(DBP) formation in distribution systems. These factors include the concentration and chemical
properties of the DBP precursors, water temperature, pH, disinfectant type, dose and residual,
and contact time (Rossman et al., 2001; Liang and Singer, 2003; Baribeau et al., 2005). Metal
pipe materials and their corrosion byproducts may also affect the formation and the fate of
halogenated DBPs in distribution systems as has been reported for iron and copper-based pipes
by Li et al. (2007), Zhang and Andrews (2012) and Rossman et al. (2001). However, no relevant
information concerning the impacts of pipe deposits on nitrosamine formation has been reported
at the time of writing, and the literature concerning the fate of nitrosamines in full-scale
distribution systems is varied (Wilczak et al., 2003; Baribeau et al., 2006; Charrois et al., 2007).
147
In addition, given the impacts of flow conditions on chloramine degradation and that chloramine
concentration is a rate-limiting factor for NDMA formation, the impacts of flow conditions on
NDMA formation merits some investigation.
Therefore, the objective of this study was to evaluate the impacts of pipe materials
(ductile iron, PVC, copper and lead) on chloramine stability and NDMA formation in the
absence and presence of orthophosphate as a corrosion inhibitor and, in particular, under
different flow conditions. These pipe materials are widely present in either water mains (such as
ductile iron and PVC) or household plumbing systems (such as copper), and lead can be
commonly found in old lead service lines, soldered joints and brass plumbing fittings. Modified
pipe loops were used in the tests under typical chloramination conditions. Results of this study
improve our understanding about key factors that affect chloramine stability and NDMA
formation in the complex physiochemical and biological environment of distribution systems.
7.2 Materials and Methods
Reagents and Materials 7.2.1
All chemicals used in this study were ACS grade or higher. Orthophosphate (K3PO4) was
the corrosion inhibitor selected for testing. The test water was the chloraminated reservoir
effluent from the Mannheim Water Treatment Plant (MWTP), Ontario. The chloramine
concentration in the test water ranged from 1.3 mg/L to 1.6 mg/L as Cl2, and other water quality
parameters in the test water are listed in Table 7-1.
Table 7-1 Summary of water quality parameters for the influent of the pipe loops
Parameters Values
pH 7.4 ± 0.1
Alkalinity (mg/L) 224±11
UV254 (cm-1
) 0.039±0.007
TOC (mg/L) 2.58 ± 0.25
SUVA (L/mgcm-1
) 0.015±0.003
Ammonia (mg/L as N) 0.17±0.05
Nitrate (mg/L as N) 4.06±0.34
148
Due to the consistently low formation of NDMA in the test water, one NDMA precursor,
dimethylamine (DMA, purchased from Sigma-Aldrich), was spiked to ensure NDMA formation
at a measurable level (> 1 ng/L). A dosing solution containing DMA and/or orthophosphate was
prepared by adding 1 mL 25 g/L (0.56 M) DMA and/or 20 g K3PO4 into 25 L unchloraminated
post-filtration water from the MWTP. The target concentrations of orthophosphate and DMA in
the influent were 1 mg/L as P and 20 µg/L (450 nM), respectively.
Modified Pipe Loops 7.2.2
Dalhousie pipe loops that were designed and manufactured by Dr. Gagnon’s research
group (Gagnon et al., 2008) were modified according to Cantor (2009) to increase their
efficiency and expand the range of operating conditions. Sixteen square test coupons (6.4 6.4
0.16 cm; Metal Samples Co., Alabama, US) were stacked inside 10 cm (ID) Schedule 40 PVC
pipes to create test modules. The module construction is shown in Figure 7-1. The main module
design parameters are summarized in Table 7-2. The lengths of the PVC-, iron-, copper- and
lead-containing modules were determined based on volume-to-surface area ratios commonly
observed in water mains and residential plumbing systems.
Figure 7-1 Module configuration (adapted from Cantor, 2009)
Flange Pipe insertion rack
4” diameter PVC pipe
Internal of module: 8 metal plates stacked on the tongue of the insertion rack, separated by plastic spacers and connected with a PVC bolt
Metal plate identification
149
Table 7-2 Summary of design parameters for pipe loops
V/S*
(mL/cm2)
Module
length (cm)
# of
plates
Equivalent distribution
system pipe diameter
(cm)
PVC 3.9 52 0 16
Iron 4.5 53 16 18
Copper 3.0 26 16 12
Lead 3.2 28 16 13
Note: * Volume-to-surface area ratio;
Volume is calculated as the sum of volumes of the return section and the test module.
The modified pipe loops are illustrated in Figure 7-2. Each pipe loop consisted of 6
components: the test module, the recirculating pump, the return section, the aluminum support
frame, the influent feed pump, and the influent and effluent ports. A PVC flexible hose was used
as the return section. The influent port was located on the return section and was the connection
point for the influent feed pump to provide fresh water to the pipe loop. The effluent port was
located on the module and acted as both a sample collection point and the outflow of the water.
The test water was introduced into the pipe loop by the influent feed pump, the flow rate of
which was adjusted based on the desired retention time and the volume of the pipe loop. The
recirculating pump propelled the water through the pipe loop to provide the desired velocity
(typically 0.3 m/s) within the test section.
Figure 7-2 Schematic of the modified pipe loop (not to scale)
Recirculation pump
Influent feed pump Influent reservoir
Return section
Transition section
Aluminum support frame DMA/PO43-
dosing tank
DM
A/P
O4
3- fe
ed
pu
mp
Effluent port
150
Experimental Procedures 7.2.3
After the pipe loops were modified, they were first flushed with the chloraminated
reservoir effluent for 3 days, and then conditioned for periods of up to 120 days to allow released
metal concentrations to stabilize in the test water matrix. During conditioning, the flow rates in
the four pipe loops were increased stepwise to obtain hydraulic retention times (HRTs) of 12
hours, 6 hours and 2 hours. To eliminate chlorine demand from the newly installed pipes, a high
concentration of free chlorine (390 mg/L) was pumped into the loops at a flow rate of 30~38
mL/min for 3 hours. Due to a persistently high chlorine demand in the copper loop, another two
conditioning treatments (180 mg/L free chlorine, 10 days) were conducted for that loop. After
these treatments, the 6 hour chloramine demand in the PVC, iron and lead loops decreased from
1.1, 1.5 and 1.5 mg/L before treatment to 0.1, 0.5 and 0.5 mg/L, respectively, and the 2 hour
chloramine demand in the copper loop decreased from 1.5 mg/L before treatment to 1.2 mg/L.
Experiments were performed in two phases: first, without the addition of orthophosphate
and then with the addition of 1 mg/L orthophosphate. In each phase, three flow regimes typically
encountered in distribution systems were examined - turbulent, laminar and stagnant. Turbulent
and laminar flow conditions were achieved by turning on and off the recirculating pumps,
resulting in flows with Reynolds numbers (Re) of approximately 30000 and <10, respectively.
HRTs in each pipe loop were varied by adjusting the influent feed pump flow rate. The stagnant
flow condition was achieved by turning off both the influent feed pump and the recirculating
pump.
Water samples from the influent and the effluents of the four pipe loops were collected
three times a week for chloramine, released metal concentrations and NDMA measurement. pH,
TOC, UV254, alkalinity, temperature, phosphate, ammonia, nitrite and nitrate were measured
once a week. To ensure a relatively constant concentration of DMA being dosed into the test
water, the flow rates of the DMA dosing solution and the reservoir effluent were monitored on
each sampling date. In addition, the DMA concentration being dosed into the influent water was
checked before and after the addition of orthophosphate by monitoring the 24 hour NDMA
formation in 1 L amber bottles with the ambient chloramine concentration.
Measurement of total organic carbon (TOC) concentration was undertaken using a Model
1030 TOC analyzer (OI Analytical, USA). UV254 was measured using a CE 3055 Reflectance
151
Spectrophotometer (Cecil Instruments, England) at 254 nm. Measurements of total chlorine and
ammonia were performed using a Hach DR2700 Spectrophotometer (Hach Company, USA) at
530 nm and 425 nm, respectively. Routinely, total chlorine was measured instead of
monochloramine because of greater temperature dependency of the monochloramine test,
however, the total chlorine residuals were shown to consist almost exclusively of
monochloramine. Metal concentrations were analyzed using a Varian SpectrAA.20 Flame
Atomic Absorption Spectrometer (Agilent Technologies, USA) after being acidified to pH <2.
Nitrite, nitrate and phosphate concentrations were measured using a Dionex DX-300 Series Ion
Chromatography System (Thermo Scientific, USA). NDMA extract preparation was carried out
according to Taguchi et al. (1994), and the instrument used for NDMA analysis was a Varian
4000 GC/MS (Agilent Technologies, USA).
7.3 Results and Discussion
Metal and Nitrogen Species Concentrations 7.3.1
Figure 7-3 compares metal concentrations in the absence and presence of orthophosphate
under two flow conditions. Generally, the released metal concentrations were lower under
laminar conditions compared with those under turbulent conditions regardless of orthophosphate
addition, which is consistent with previous studies in that high flow rates can loosen the
corrosion byproducts deposited on the pipe wall and cause more metal to be released from the
pipe surface (Mutoti et al., 2007).
With respect to specific impacts of orthophosphate, as shown in Figure 7-3, the results
varied with both the type of pipe materials employed and the hydraulic conditions present. The
addition of orthophosphate significantly reduced the released copper concentrations under both
laminar and turbulent conditions, as expected (Edwards et al., 2002). Also, as expected,
orthophosphate did not exhibit obviously beneficial effects on iron corrosion control under both
laminar and turbulent conditions, as also evidenced by McNeill and Edwards (2000; 2001).
However, orthophosphate effectively reduced lead release but mainly when water was turbulent.
In the absence of orthophosphate, the released lead concentrations varied from 0.4 to 0.6 mg/L,
but they were below 0.1 mg/L after the addition of orthophosphate. Under laminar conditions,
152
there were no significant differences in lead concentrations before and after the addition of
orthophosphate (varying between 0.1 and 0.2 mg/L).
Figure 7-3 Metal concentrations in the pipes of iron, copper and lead in the absence and presence
of 1 mg/L orthophosphate under two flow conditions
Given that orthophosphate previously has been shown to effectively decrease lead release
(Edwards et al., 1999; Edwards and McNeill, 2002), the observed relative ineffectiveness of
orthophosphate on lead corrosion control under laminar conditions in the current study might be
explained by the relative newness of the module materials (leading to higher than normal lead
release rates) and/or interplays among lead and dissolved species such as ammonia, nitrite and
0.0
0.2
0.4
0.6
0.8
0 3 6 9 12 15
Pb
(m
g/L
)
HRT (hours)
No phosphate spiked
Phosphate spiked
0.0
0.2
0.4
0.6
0.8
0 3 6 9 12 15
Pb
(m
g/L
)
HRT (hours)
No phosphate spiked
Phosphate spiked
0.0
0.2
0.4
0.6
0.8
1.0
0 3 6 9 12 15
Fe
(m
g/L
)
HRT (hours)
No phosphate spiked
Phosphate spiked
0.0
0.2
0.4
0.6
0.8
1.0
0 3 6 9 12 15
Fe
(m
g/L
)
HRT (hours)
No phosphate spiked
Phosphate spiked
0.0
0.5
1.0
1.5
2.0
0 3 6 9 12 15
Cu
(m
g/L
)
HRT (hours)
TURBULENT
No phosphate spiked
Phosphate spiked
0.0
0.5
1.0
1.5
2.0
0 3 6 9 12 15
Cu
(m
g/L
)
HRT (hours)
LAMINAR
No phosphate spiked
Phosphate spiked
153
nitrate (Edwards and Dudi, 2004; Uchida and Okuwaki, 1998; Zhang et al., 2009a) as shown in
Reactions 7-1 and 7-2:
NO3- + Pb NO2
- + PbO Reaction 7-1
NO2- + 3Pb +2H2ONH3 + 3PbO+OH
- Reaction 7-2
According to these reactions, free ammonia from the application of monochloramine combined
with a high concentration of nitrate may have synergistically increased the lead corrosion by
interfering with the formation of the passive PbO layer on its surface. If this phenomenon was
occurring, then according to Reactions 7-1 and 7-2, there should be a strong correlation among
the concentrations of the three nitrogen species. Namely, a decrease in nitrate concentrations
should lead to an increase in both nitrite and ammonia concentrations. Figure 7-4 illustrates the
variations in the nitrogen species concentrations measured in the lead loop. It confirms that
decreases in nitrate concentrations were associated with increases in both nitrite and ammonia
concentrations and provides evidence for the potential occurrence of Reactions 7-1 and 7-2 in the
lead loop. The co-presence of ammonia, nitrite and nitrate in significant concentrations in the
lead loop may also suggest the occurrence of nitrification as a potential source of these ions, but
confirmation of this was beyond the scope of the study and requires further investigation.
Figure 7-4 Variations of nitrate, nitrite and ammonia in the lead loop
0.0
0.1
0.2
0.3
3.0
3.5
4.0
4.5
5.0
27-Apr-11 17-May-11 6-Jun-11 26-Jun-11
Am
mo
nia
/nitrite
(m
g/L
-N)
Nitra
te (
mg
/L-N
Nitrate Nitrite Ammonia
154
Chloramine decay 7.3.2
Generally, chloramine concentrations in the four pipe loops all followed pseudo-first-
order decay kinetics, so decay constants were estimated by fitting a pseudo-first-order decay
equation to the chloramine data using Microsoft Excel.
7.3.2.1 Effects of Orthophosphate and Flow Conditions
Figure 7-5 displays pseudo-first-order chloramine decay constants for the four pipe
materials before and after the addition of orthophosphate under three flow regimes (Reynolds
numbers of Re=30000, Re 10, and Re=0). In addition, for each pipe material, a Fisher’s Least
Significant Difference (LSD) test at a confidence level of 95% was applied to pairs of
chloramine decay constants that were obtained with different treatments to estimate the
significance of the effects of different flow conditions and the presence or absence of
orthophosphate on chloramine decay (Montgomery, 2000). The results of the LSD tests to
determine the effects of flow conditions are summarized in Table 7-3.
Figure 7-5 Pseudo-first-order chloramine decay constants for the four pipe loops under different
flow conditions, initial chloramine 1.3~1.6 mg/L as Cl2, error bars indicate the measured
maximum and minimum values (n=2)
0.0
0.2
0.4
0.6
0.8
1.0
Re=0 Re<10 Re=30000
De
ca
y c
on
sta
nt (h
-1)
Fe
0.0
0.2
0.4
0.6
0.8
1.0
Re=0 Re<10 Re=30000
Deca
y c
on
sta
nt (h
-1)
Cu
0.0
0.2
0.4
0.6
0.8
1.0
Re=0 Re<10 Re=30000
Deca
y c
on
sta
nt (h
-1)
Pb
0.0
0.2
0.4
0.6
0.8
1.0
Re=0 Re<10 Re=30000
Deca
y c
on
sta
nt (h
-1)
PVC
No orthophosphate Orthophosphate spiked
155
Table 7-3 Significance of the effects of flow conditions on chloramine decay determined by the
LSD test (95% confidence level)
Treatments compared Fe Pb Cu PVC
No
orthophosphate
Re=0 Re<10 ↓
Re<10 Re=30000 ↑ ↓ ↑ ↑
Re=0 Re=30000 ↑ ↑ ↑ ↑
Orthophosphate
spiked
Re=0 Re<10 ↑ ↓ ↓ ↑
Re<10 Re=30000 ↑ ↑ ↑
Re=0 Re=30000 ↑ ↑ ↑
Note: , no significant effect; ↑, increasing decay; ↓, decreasing decay
Generally, chloramine decay constants for all four pipe materials increased with
increasing water flow. This agrees with previous findings that an increase in flow velocity
increases the mass transfer rate of chloramine to the pipe surface and/or increases the release rate
of metal corrosion products from the pipe surface, thereby increasing chloramine consumption
(Rossman et al., 1994; Westbrook and Digiano, 2009). The results for the PVC loop also agree
with a previous finding that bulk chlorine decay rates are dependent on flow velocity (Menaia et
al., 2003).
The decay constants were also influenced by the type of pipe material present, which is
newly identified in this study. As shown in Figure 7-5, there was a marked increase in
chloramine decay in both the lead and copper pipes under turbulent conditions. The rapid loss of
chloramine was observed to occur in the initial 1~2 hours of the tests during which time
approximately 70% of chloramine was degraded.
In the lead loop, the chloramine decay may have been enhanced by the presence of nitrate
in the water. As discussed in Section 7.3.1, the capability of lead to react with nitrate has been
well documented (Uchida and Okuwaki, 1998; Edwards and Dudi, 2004; Zhang et al., 2009a).
Reactions 7-1 and 7-2 can essentially increase nitrite formation and lead corrosion rates. Under
turbulent conditions, up to 0.08 mg/L nitrite as N and 0.4 mg/L Pb (Figures 7-3 and 7-4) were
detected in the lead pipe in the absence of orthophosphate, and they were 40% and 75% higher,
respectively, than those in the presence of orthophosphate. Since chloramine reacts with NO2-
and considering that NO2- is formed in Reaction 7-1, it was reasonable to have observed
accelerated chloramine degradation in the lead loop.
156
The high chloramine demand in the copper pipe under turbulent conditions is consistent
with previous findings that chloramine can be rapidly consumed by copper corrosion via the
following reaction (Nguyen, 2005):
½NH2Cl +H+ +½Cu
0 + ½Cu
2+ Cu
2+ + ½NH4
+ + ½Cl
- Reaction 7-3
Under turbulent conditions, the copper surface can be scoured, exposing fresh copper and
providing a virtually unlimited supply of metallic copper (Cu0) to participate in this reaction. The
observed rapid chloramine decay under these conditions is, therefore, primarily due to the high
chlorine demand from the direct oxidation of fresh copper to cupric species.
To evaluate the significance of the presence or absence of orthophosphate on chloramine
decay, the LSD tests were also applied and the results are summarized in Table 7-4. As shown in
Table 7-4 and also in Figure 7-5, the addition of orthophosphate did not cause statistical changes
in chloramine degradation at Re =0 and Re <10 for both the copper and PVC loops. However,
under turbulent flow (Re =30,000), orthophosphate decreased chloramine degradation in the
copper loop, and the opposite effect was observed in the PVC loop. Surprisingly, orthophosphate
significantly increased chloramine decay in the iron loop, but only at Re =0 and Re <10 for the
lead loop.
Table 7-4 Significance of the effects of orthophosphate on chloramine decay determined
by the LSD test (a confidence level of 95%)
Materials Treatments Re=0 Re<10 Re=30000
Fe
No
orthophosphate
Orthophosphate
spiked
↑ ↑ ↑
Pb ↑ ↑ ↓
Cu ↓
PVC ↑
Note: , no significant effect; ↑, increasing decay; ↓, decreasing decay
For conditions under which orthophosphate significantly increased chloramine decay (in
the iron and lead loops), Figure 7-3 indicates that metal release kinetics were comparable before
and after the addition of orthophosphate. This suggests that, rather than metal corrosion
processes, other factors (e.g. nitrite formation) may have played an important role in accelerating
chloramine decay. Further discussions about this theory are provided in Section 7.3.2.2.
157
7.3.2.2 Chloramine Decay and Nitrite Formation
Nitrite (NO2-) may be present from either or both of abiotic reactions (e.g. Reaction 7-1
when in contact with lead) and from nitrification (Wilczak et al., 1996), and its impacts on
chloramine decay for each type of pipe material present were further investigated in this study.
Although the presence of nitrifiers could not be detected directly in this study, nitrite
concentrations were measured in the effluents of the four pipe loops. These data are plotted in
Figure 7-6 along with the corresponding chloramine pseudo-first-order decay constants for two
tested flow conditions (turbulent and laminar).
Figure 7-6 Box and Whisker plots for nitrite concentrations before and after the addition of
orthophosphate overlaid with pseudo-first-order chloramine decay rate constants, laminar flow
(n=5), turbulent flow (n=6). The top and bottom of the box represent the 75th and 25th
percentile, respectively, while the whiskers represent the maximum and minimum values
Nitrite Decay constants
0.0
0.2
0.4
0.6
0.8
1.0
0.0
0.1
0.2
0.3
Pse
ud
o-f
irst-
ord
er
de
ca
y
co
nsta
nt (h
-1)
Nitri
te (
mg
/La
s N
) LAMINAR
0.0
0.2
0.4
0.6
0.8
1.0
0.0
0.1
0.2
0.3
Pse
ud
o-f
irst-
ord
er
de
ca
y
co
nsta
nt (h
-1)
Nitri
te (
mg
/L a
s N
) TURBULENT
158
In general, a high concentration of nitrite corresponded to a rapid chloramine degradation
constant under both laminar and turbulent conditions, as expected. In addition, under laminar
conditions, the iron and lead loops exposed to orthophosphate had a relatively higher
concentration of nitrite (~ 0.2 mg/L) compared with other pipe materials (mostly below 0.05
mg/L). This was in agreement with previous findings that orthophosphate can enhance
nitrification and, therefore, nitrite formation (Zhang and Edwards, 2010b). However, the
measured nitrite concentration did not always increase with the addition of orthophosphate, as
observed for the iron and lead loops under turbulent flow conditions. Turbulent flow can disturb
the stability of biofilms where nitrifiers could be present, thereby decreasing their activity
(Zhang et al., 2009b). Decreased nitrite formation in the lead loop after the addition of
orthophosphate under turbulent conditions is also likely because the protective layer on the lead
surface decreased the exposure of lead to nitrate in the water (Reaction 7-1).
Nitrite concentrations in the copper loop were consistently lower than in the other loops,
with the maximum concentrations for both flow conditions being below 0.02 mg/L (Figure 7-6).
It is known that high levels of copper (> 0.1 mg/L) are toxic to nitrifiers, thereby inhibiting
nitrification (Zhang et al., 2009b) and the measured copper concentrations in this study ranged
from 0.08 to 1.0 mg/L (Figure 7-3). The lack of nitrite in the copper loop also suggests that
abiotic consumption of chloramine due to copper corrosion (Reaction 7-3) was a likely reason
for the rapid chloramine dissipation under turbulent flow (Figure 7-5).
NDMA formation 7.3.3
In addition to determining the potential effects of pipe materials, orthophosphate and
flow conditions on NDMA formation, the pipe loops were operated at different hydraulic
retention times (HRTs) to evaluate the impacts of reaction time and provide additional
confirmation of the observed results.
7.3.3.1 Effects of Flow Conditions
Figure 7-7 provides a general overview of the results by comparing NDMA
concentrations in the effluent of the four pipe loops in the absence of orthophosphate under
turbulent (Re =30000) and laminar (Re <10) conditions. Results for tests performed with the
addition of 1 mg/L orthophosphate were similar, except that all of the NDMA concentrations
159
were reduced (Appendices Figure 10-21, to be discussed further in later section). Generally,
NDMA formation in each pipe loop increased with increasing HRT, as expected. Under
turbulent conditions, the sequence of NDMA formation potentials in the four pipe loops followed
the order of Cu Pb< Fe< PVC, whereas chloramine decay constants followed the opposite order
(Figure 7-5), which is in agreement with common knowledge that chloramine presence is a rate-
limiting factor for NDMA formation.
Figure 7-7 NDMA formation in four pipe loops in the absence of orthophosphate under two flow
conditions
It was interesting to observe that changing flow conditions changed these relationships.
NDMA concentrations were generally higher under laminar conditions than turbulent conditions
and the pipe materials exhibited even more pronounced but different effects on NDMA
formation. NDMA formation did not always directly correlate with chloramine presence. For
0
40
80
120
160
200
2 6 12
ND
MA
(ng/L
)
HRT (hours)
TURBULENT
0
40
80
120
160
200
2 6 12
ND
MA
(ng/L
)
HRT (hours)
LAMINAR
160
example, although chloramine decay constants in the four pipe loops were statistically similar
under this condition (Figure 7-5), the NDMA formation potentials followed a distinctly different
sequence (an increasing order of Fe, Pb, PVC and Cu, as illustrated in Figure 7-7). In particular,
the measured NDMA concentration in the copper loop tended to plateau at approximately 180
ng/L after 6 hours, but in the PVC loop (the control loop) there was only 65 ng/L NDMA at HRT
of 6 hours and 150 ng/L NDMA present after 12 hours. Thus, the chloramine decay constants
were comparable but NDMA formation was enhanced in the copper loop relative to the PVC
loop, indicating that Cu(II) may have catalyzed NDMA formation under laminar flow conditions.
In addition, for the iron loop, although chloramine degraded more quickly under turbulent
conditions than under laminar flow (Figure 7-5 and Table 7-3), Figure 7-7 shows that NDMA
yields after HRT 6 hours under turbulent conditions were at least 70% higher than those under
laminar conditions. Catalytic effects of iron corrosion products during NDMA formation may,
therefore, be suggested. In general, reactions mediated by surface-bound metal are affected by
the mineral’s surface area and the density of sorbed metal ions (Chun et al., 2005). Under
turbulent conditions, the high velocity of the water may enhance the release of iron corrosion
products from the pipe wall, thereby increasing the surface area of iron corrosion products in
contact with DMA. As a consequence, NDMA formation was likely catalyzed by iron corrosion
products under turbulent flow in the current tests despite the relatively rapid chloramine
degradation. These observations concerning the dependency of copper and iron influences on
flow conditions have not been reported previously.
7.3.3.2 Effects of Orthophosphate
Since orthophosphate’s corrosion preventive effects can generally reduce the impacts of
the tested parameters on corrosion rates and metal concentrations, this study also investigated the
effects of orthophosphate on NDMA formation. However, partway through the experimental
plan, the treatment plant performed some operational changes (increasing polyaluminum
chloride, PACl) to improve their filter performance. Coincidently, NDMA yields in the four pipe
loops significantly decreased. Since the make-up and flow rate of the orthophosphate/DMA
dosing solution and the flow rate of the reservoir effluent remained constant before and after the
operational changes, the decreased NDMA formation was most likely attributed to the change in
quality of the chloraminated reservoir water as a result of the increased PACl addition. Such
161
effects of water matrix on NDMA formation has also been reported by Oya et al. (2008).
Regardless, to account for variations in water quality, the effects of orthophosphate on NDMA
formation were determined after first normalizing the NDMA concentrations in the different
loops relative to those observed in the PVC control loop.
Normalized NDMA concentrations are shown in Figure 7-9. Since the calculated relative
percentage values of NDMA formation in the PVC loop were all 100%, they are not shown
individually, but are indicated by the horizontal line at 100%.
Figure 7-9 NDMA formation relative to the PVC control in the absence and presence of 1 mg/L
orthophosphate under laminar and turbulent flow conditions
Under turbulent conditions, NDMA formation relative to the PVC loop in all of the three
metal loops in the presence of orthophosphate increased compared with that in the absence of
orthophosphate. The greatest increase was observed in the iron loop, in which more NDMA
formed than in the PVC loop after the addition of orthophosphate. Recall that there was higher
0
50
100
150
200
250
300
350
2 6 12 2 6 12
ND
MA
form
ation r
ela
tive t
o the
PV
C c
ontr
ol (%
)
TURBULENT
Fe Cu Pb
No orthophosphate Orthophosphate
0
50
100
150
200
250
300
350
2 6 12 2 6 12
ND
MA
form
ation r
ela
tive t
o the
PV
C c
ontr
ol (%
)
HRT (hours)
LAMINAR
Fe Cu Pb
No orthophosphate Orthophosphate
162
NDMA formation but more rapid chloramine decay in the iron loop under turbulent conditions
than laminar conditions, suggesting the possible catalysis of iron on NDMA formation (Section
7.3.3.1). The additional evidence of increased NDMA formation after the addition of
orthophosphate in the iron loop indicates that orthophosphate can significantly increase these
catalytic effects of iron corrosion products, possibly by affecting the properties of iron particles
and their suspension in the solution (Lytle and Snoeyink, 2002). As for the copper and lead
loops, according to Table 7-4, orthophosphate significantly decreased chloramine decay by
reducing metal corrosion rates. Since chloramine decay is a rate-limiting factor for NDMA
formation, it is also reasonable to have observed the increased formation of NDMA in these two
loops after the addition of orthophosphate.
Under laminar conditions, the main observation is that NDMA formation relative to the
PVC control in the copper loop is significantly higher than 100%, especially in the absence of
orthophosphate. This confirms the trends shown in Figure 7-7 that illustrate copper catalysis of
NDMA formation under laminar conditions. Furthermore, the addition of orthophosphate
decreased the percentage values of NDMA formation relative to the PVC control in the copper
loop for all the HRTs. This suggests orthophosphate may mitigate copper catalysis on NDMA
formation by decreasing the released copper concentrations (Figure 7-3).
7.4 Summary
This is the first study to employ modified pipe loops to determine the effects of pipe
materials, flow conditions and orthophosphate on chloramine decay and NDMA formation from
DMA. Ductile iron, copper, lead and PVC were tested. Three flow regimes encountered in
distribution systems (turbulent, laminar and stagnant) were examined. In general, the pipe
materials, flow conditions and orthophosphate were all observed to influence metal corrosion,
chloramine decay and NDMA formation as follows:
a. In agreement with previous studies, turbulent conditions generally increased the released
metal concentrations compared with laminar conditions regardless of orthophosphate
addition. Orthophosphate had no beneficial effects on iron corrosion control, but
effectively reduced copper concentrations under both laminar and turbulent conditions.
163
b. Chloramine degradation kinetics for four pipe materials were compared. The new
observation was that the impacts of flow conditions and orthophosphate on chloramine
decay were highly dependent on the type of pipe material. Orthophosphate increased
chloramine degradation in the iron loop regardless of flow conditions. Abiotic copper
corrosion was the primary reason for the accelerated chloramine decay under turbulent
condition. For the lead loop, a rapid dissipation of chloramine was primarily due to the
electrochemical reactions between metallic lead with nitrate to form nitrite. The addition
of orthophosphate mitigated the impacts of turbulence on chloramine decay in both the
copper and lead loops.
c. Consistent with previous studies, chloramine concentration was generally a rate-limiting
factor for NDMA formation, especially under turbulent conditions.
d. The impacts of flow conditions on metal catalysis of NDMA formation were newly
identified. Regardless of the presence of orthophosphate, copper catalyzed NDMA
formation from DMA under laminar conditions, but iron catalyzed NDMA formation
only under turbulent conditions. Iron catalysis was increased by the addition of
orthophosphate likely because orthophosphate modified properties of the associated
suspended particles, while orthophosphate decreased copper catalysis likely by reducing
dissolved copper concentrations.
Results of this study help in examining the complex reactions involving metal corrosion,
chloramine degradation and NDMA formation under different flow conditions in distribution
systems. Since it is an initial study to have identified the catalytic impacts of iron and copper on
NDMA formation and their dependency on flow conditions, further study is recommended to
investigate the mechanisms for iron and copper catalysis. Nevertheless, knowing more about
impacts of pipe materials and corrosion inhibitors as well as the possible occurrence of
nitrification on the stability of chloramine residuals and subsequent NDMA formation will be
useful in developing strategies to control metal corrosion and reduce disinfection by-product
formation in distribution systems.
164
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haloacetic acids and trihalomethanes in drinking water. Environmental Science &
Technology, 37(13), 2920-2928.
Lytle, D. A., and Snoeyink, V. L. (2002) Effect of ortho- and polyphosphates on the properties of
iron particles and suspensions. Journal American Water Works Association, 94(10), 87-
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McNeill, L. S., and Edwards, M. (2000) Phosphate inhibitors and red water in stagnant iron
pipes. Journal of Environmental Engineering-Asce, 126(12), 1096-1102.
McNeill, L. S., and Edwards, M. (2001) Iron pipe corrosion in distribution systems. American
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168
Conclusions 8
The specific objectives of this thesis outlined in Section 1.2 were addressed in this work and
resulted in some main conclusions. Generally, of the results discussed below, those that are
summarized in Conclusion 1 (from Chapter 3) compared the impacts of corrosion inhibitors and
their interactions with metal surfaces on chlorine degradation and HAA9 formation, and the
results in Conclusions 3 to 5 (from Chapters 5 to 7) in terms of copper catalysis and flow
conditions on HAA/NDMA formation are essentially new contributions to the field. The results
in Conclusion 2 (from Chapter 4) concerning the impacts of disinfectant type, corrosion
inhibitors and water quality on metal release kinetics confirmed those of previous studies.
The main conclusions of this thesis include:
1) Phosphate-based corrosion inhibitors and their interactions with metal materials affected
the kinetics of disinfectant degradation (both HOCl and NH2Cl) and DBP formation
(mainly HAAs) in two types of water matrices, depending on metal type and metal age.
Newly discovered was that HAA formation was enhanced in the presence of high levels of
copper ions from fresh metal coupons in both investigated water matrices, indicating
possible catalytic potential of copper on HAA formation. The addition of phosphate-based
corrosion inhibitors alone generally did not impact HAA formation in both tested water
matrices. Enhanced formation of NDMA from DMA in water containing fresh iron and
copper coupons also indicated the possible catalysis of iron and copper on NDMA
formation, and that the catalysis may increase with the released metal concentrations.
Consistent with previous studies, both HOCl and NH2Cl generally followed pseudo-
first-order kinetics regardless of metal age and water quality. The reactivity of these pipe
materials with HOCl and NH2Cl followed the sequence of decreasing order of ductile iron,
copper and lead. For fresh metal coupons in both tested water matrices, the addition of
phosphate-based corrosion inhibitors significantly increased chlorine decay for iron
coupons, but reduced chlorine decay for copper coupons. For corroded coupons in both
tested water matrices, however, orthophosphate did not statistically affect free chlorine
decay, regardless of metal type.
2) Disinfectant types, corrosion inhibitors and water quality affected metal release kinetics in
simulated distribution systems. Early experiments generally confirmed the results of other
169
researchers and also suggested that orthophosphate increased the level of released iron,
regardless of the age of metal surface, water quality, and disinfectant type. However,
orthophosphate significantly decreased released copper concentrations, in particular for
short term exposures to fresh copper metal. For corroded copper coupons, an increased
copper corrosion rate was observed in the presence of orthophosphate under chlorination
for the water with high pH and alkalinity values. Lead release was significantly reduced in
the presence of orthophosphate, irrespective of the age of metal surface and water quality.
Regardless of orthophosphate addition, high levels of metal ions were released in the water
with low pH and alkalinity, which is consistent with other research.
3) Copper corrosion products, including Cu(II), Cu2O, CuO, and Cu2(OH)2CO3, exhibited
catalytic effects on HOCl decay and HAA formation. Copper catalysis was affected by pH
and the concentration of these corrosion products. The presence of Cu(II) and its solid
corrosion products led to DCAA formation consistently predominating over the formation
of other HAA species. Further investigation of chlorine decay pathways in the presence of
Cu(II) in synthetic water indicated that Cu(II) would interact with NOM, possibly by
complexation, and increase the reactivity of NOM with chlorine. As a result, chlorine
decay was accelerated by reacting with these active Cu(II)-NOM complexes and HAA
formation was enhanced.
4) Copper was shown to catalyze NDMA formation from DMA. NDMA formation from DMA
was increased with increasing Cu(II) concentrations, DMA concentrations, alkalinity and
hardness, but was inhibited by the presence of NOM. The rapid consumption of NH2Cl by
NOM and/or the competitive complexation of NOM with copper were proposed to be
involved in limiting NDMA formation by NOM. pH influenced the speciation of
chloramine and the interactions of copper with DMA. Elevated formation of NDMA at
neutral pH was primarily attributed to the transformation of monochloramine to
dichloramine and complexation of copper with DMA. In addition, aqueous copper released
from malachite [Cu2CO3(OH)2] was shown to promote NDMA formation while the
presence of CuO decreased NDMA formation.
5) Turbulent conditions were shown to increase chloramine decay, but affected NDMA
formation differently, particularly for copper and iron. Orthophosphate increased
chloramine degradation in the iron loop irrespective of flow conditions. Accelerated
chloramine decay was observed in the copper loop under turbulent flow conditions
170
primarily due to abiotic copper corrosion. The addition of orthophosphate effectively
decreased chloramine decay by reducing the released copper concentrations. Rapid
chloramine decay was also observed in the lead loop under turbulent conditions, but the
addition of orthophosphate mitigated the impacts of turbulence on chloramine decay. Bulk
water chlorine reactions in the PVC loop were also increased with increasing flow velocity,
as has also been reported by others.
Chloramine concentration was generally a rate-limiting factor for NDMA
formation, and high chlorine demands due to metal corrosion and/or nitrite formation was
associated with reduced NDMA formation. Pipe materials affected the transformation of
NDMA to DMA: copper consistently exhibited its catalysis on NDMA formation from
DMA under laminar conditions, whereas iron catalyzed NDMA formation only under
turbulent conditions. Orthophosphate was shown to reduce catalytic effects of copper but
appeared to increase iron catalysis by affecting the metal corrosion processes.
Practical Implications and Suggestions for Future Research 9
This research investigated the impacts of metal corrosion and corrosion inhibitors on
disinfectant residual degradation and DBP formation (HAA and NDMA) at both bench- and
pilot-scale. Metal materials that were investigated include ductile iron, copper and lead. Key
factors affecting metal corrosion, disinfectant residual degradation and subsequent DBP
formation were examined, including metal age, water quality, and flow conditions. The catalytic
potential of copper to affect HAA and NDMA formation were evaluated under controlled
experimental conditions at bench-scale, with further experiments concerning NDMA formation
being performed at pilot-scale.
The findings about copper catalysis during HAA and NDMA formation provide
important implications for distributed water quality in domestic plumbing systems where
primarily copper pipes are installed. These pipe surfaces are often covered by corrosion solids.
Interactions of disinfectant residuals and DBP precursors with copper corrosion products will
affect the stability of secondary disinfectants and the fate of DBPs. Understanding the catalytic
potential of copper on disinfectant degradation and HAA/NDMA formation will be of benefit to
water utilities and households for the management of their distribution systems and water
quality. Since copper concentration, pH, reaction time and flow conditions are the main factors
171
impacting the nature and extent of copper catalysis in the present study, considerations for
corrosion preventive strategies are recommended to include: pH adjustment and addition of
corrosion inhibitors at the treatment plant to control copper corrosion, and flushing prior to use in
the home to decrease contact time of water with pipe corrosion products. These steps will help to
alleviate HAA and NDMA formation by copper catalysis and ensure safe and clean water is
delivered to the consumer.
This research helps achieve a better understanding of the complex reactions involving
metal corrosion, disinfectant residual degradation and DBP formation. The results provide
insight into the factors affecting disinfectant residual stability and the fate of DBPs in
distribution system water mains and household plumbing. Although phosphate-based corrosion
inhibitors have received increasing attention as an effective strategy for corrosion control, some
of these corrosion inhibitors have been shown to exhibit either beneficial or detrimental effects
on disinfectant decay and subsequent DBP formation. The effects observed in the present
research have depended on the metal types, metal age, hydrodynamic conditions and water
quality. As well, these potential impacts should be considered site-specific.
In general, results from this research will provide water utilities with insight into
efficiently targeting their water characteristics and coordinating, spatially and temporally, the
critical factors to maintain their distributed water quality. For example, booster chlorination or
chloramination facilities may be used for the pipe systems with high water age or where rapid
disinfectant degradation occurs (e.g. in turbulent conditions). Controlling water chemistry (such
as pH and alkalinity) has been confirmed to be able to minimize metal corrosion, and thus it can
be considered when the addition of phosphate-based corrosion inhibitors is not effective for
corrosion control (e.g. iron). Selection of alternative pipe materials to control metal leaching will
also mitigate aesthetic and public health impacts caused by increased metal concentrations.
Limitations of this research and recommendations for future study include:
1) The impacts of copper catalysis on NDMA formation need to be investigated further in
different water matrices. Although bench-scale experiments were performed to investigate
factors of alkalinity, pH, NOM, and hardness affecting copper catalysis during NDMA
formation from DMA in Milli-Q water under controlled experimental conditions (Chapter
6), in real water, different combinations of levels of these parameters may complicate
172
NDMA formation. For instance, different pH and alkalinity conditions may affect the
speciation and solubility of hardness ions. As well, NOM in different water matrices has
different compositions and binding affinities with copper, which may affect the interactions
between DMA and copper.
2) Both bench- and pilot-scale experiments involving NDMA formation employed DMA as a
model or surrogate NDMA precursor (Chapters 6 and 7). This particular precursor may
only be relevant in worst-case scenarios in distribution systems where wastewater
significantly impacts drinking water sources or intrusion occurs introducing a significant
amount of DMA into the distributed water. However, it is unlikely that there are significant
concentrations of DMA in drinking water. Factors affecting NDMA formation from other
precursors in distribution systems needs to be examined. These precursors may include
tertiary amines containing DMA moieties, amine containing polymers, and humic
substances typically present in drinking water.
3) Pilot-scale pipe loop experiments investigated the impact of flow conditions on total
chlorine decay and NDMA formation (Chapter 7), however, only three flow conditions
were considered (Re=0, Re<10 and Re=30000 representing stagnant, laminar and turbulent
flow in distribution systems, respectively). A wider range of Reynolds numbers commonly
encountered in distribution systems should be studied in the future for a better
understanding of the fate of NDMA in distribution systems.
4) In pilot-scale pipe loop experiments, the enhanced formation of NDMA from DMA by iron
was hypothesized to be due to iron catalysis (Chapter 7), however, further work with iron
could not be accommodated in this thesis. Therefore, the mechanisms related to catalysis of
iron corrosion products on NDMA formation need to be identified and the critical factors
influencing the catalytic potential of these corrosion products need to be evaluated.
173
Appendices 10
10.1 Authorisations to Include Copyright Material in Thesis
NRC RESEARCH PRESS LICENSE
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Mar 09, 2012
This is a License Agreement between Hong Zhang ("You") and NRC Research Press ("NRC
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All payments must be made in full to CCC. For payment instructions, please see information
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License Number 2864841496472
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Licensed content publisher NRC Research Press
Licensed content publication Canadian Journal of Civil Engineering
Licensed content title Effects of phosphate-based corrosion inhibitors on the kinetics of chlorine
degradation and haloacetic acid formation in contact with three metal materials
Licensed content author Hong Zhang et al.
Licensed content date Jan 1, 2012
Volume number 39
Issue number 1
Type of Use Thesis/Dissertation
Requestor type Author (original work)
Format Electronic
174
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EFFECTS OF DISINFECTANT CHANGES AND CORROSION CONTROL STRATEGIES
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ELSEVIER
Ways to Use Journal Articles Published by Elsevier: A Practical Guide
How authors can reuse their own articles published by Elsevier
General use of articles
May 18, 2012
Authors publishing in Elsevier journals retain wide rights to continue to use their works to
support scientific advancement, teaching and scholarly communication.
An author can, without asking permission, do the following after publication of the author’s
article in an Elsevier-published journal:
• Make copies (print or electronic) of the article for personal use or the author’s own classroom
teaching
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colleagues for their personal use but not for commercial purposes or systematic distribution as
defined on page 3 of this pamphlet
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• Allow the author’s employer to use the article in full or in part for other intracompany use (e.g.,
training)
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article
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writings and lecture notes
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177
10.2 QA/QC Protocols
Various QA/QC measures were undertaken in this research to ensure analytical precision
and accuracy.
1) High purity organic solvents (HPLC grade) were used in analyses. All other chemicals
were at least of analytical grade.
2) All glassware used in experiments were made chlorine demand free before use by
soaking them in a concentrated sodium hypochlorite solution (~1000 mg/L as Cl2) for at
least 24 hours. Thereafter, the bottles were rinsed thoroughly with deionized water and
distilled water, and were heated at 250 °C for at least 4 hours.
3) Surrogate standards were used for HAA9 and nitrosamine analysis to correct any errors
from the extraction and instrumental analysis.
4) Calibration curves for each species of HAA9, NDMA, metals and anions were established
in the range representative of actual sample concentrations. Calibration standards were
prepared and run with each set of samples. Calibration curves for NDMA, HAA9, metals
and anions are provided in Figures 10-1, 10-3, 10-4, 10-11, 10-12 and 10-13.
5) At the beginning of each analytical run, lab water blanks and sample blanks containing
the surrogate standard were injected to condition the instrument and to verify that
interferences were absent.
6) Spike and recovery standards were processed and run with each set of samples to validate
and assess the accuracy of instrumental analysis. Recoveries should range between 70%
and 130% (USEPA, 2003 and 2004). Spike recovery charts for NDMA, MCAA, MBAA,
DCAA, TCAA, DBAA and BCAA are illustrated in Figures 10-2 and 10-5~10-10.
7) Method detection limits (MDLs) of the instrumental methods of analysis for each organic
compound, metal ion and anion of interest were determined by multiplying the standard
deviation of 8 replicates by the Student t value, and are provided in Tables 10-1~10-4.
8) All of the bench-scale and pilot-scale kinetic experiments were conducted, at a minimum,
in two sets with duplicate testing involved in each set in order to test the reproducibility
of the results and to determine the errors associated with sampling, extraction and
measurement for metal ions, disinfectant residual, HAA9, and nitrosamines.
178
GC/MS 10.2.1
Figure 10-1 Example of GC/MS calibration curve for NDMA (NDMA 0~200 ng/L, d6-NDMA
50 ng/L)
Table 10-1 GC/MS method – Method detection limits for NDMA (n=8, 99% confidence level)
Compound Spiked
Concentration
(ng/L)
%
Recovery RSD (%)
MDL
(ng/L)
NDMA 1 ng/L 93 31.8 0.88
y = 55.153x - 5.543 R² = 0.9993
0
50
100
150
200
250
0.0 1.0 2.0 3.0 4.0
ND
MA
co
nce
ntr
atio
n (
ng/L
)
Ratio of peak areas (NDMA/d6-NDMA)
179
Figure 10-2 GC/MS method for NDMA- spike recovery chart
0
20
40
60
80
100
120
140
0 10 20 30 40 50 60
Sp
ike
re
cove
ry (
%)
Sample set
2010 2011
130%
70%
180
GC/ECD 10.2.2
Figure 10-3 Example of GC-ECD calibration curves for HAA9 (HAA9 0~100 µg/L, 2,3,5,6-
TFBA 100 µg/L)
y = 28.898x + 1.4759 R² = 0.9932
0
20
40
60
80
100
120
0.0 1.0 2.0 3.0 4.0
Concentr
ations (
µg/L
)
Ratio of peak areas
DCAA
y = 10.783x + 0.3458 R² = 0.9967
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0 10.0
Concentr
ations (
µg/L
)
Ratio of peak areas
TCAA
y = 283.7x + 2.3356 R² = 0.9881
0
20
40
60
80
100
120
0.0 0.1 0.2 0.3 0.4
Concentr
ations (
µg/L
)
Ratio of peak areas
MCAA
y = 33.09x + 2.9221 R² = 0.9911
0
20
40
60
80
100
120
0.0 1.0 2.0 3.0 4.0
Concentr
ations (
µg/L
)
Ratio of peak areas
MBAA
y = 13.485x + 1.7596 R² = 0.9964
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0
Concentr
ations (
µg/L
)
Ratio of peak areas
BCAA
y = 12.571x + 0.2658 R² = 0.9984
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0 10.0
Concentr
ations (
µg/L
)
Ratio of peak areas
DBAA
181
Figure 10-4 Example of GC-ECD calibration curves for HAA9 (HAA9 0~100 µg/L, 2,3,5,6-
TFBA 100 µg/L), continued
Table 10-2 GC-ECD method – Method detection limits for HAA9 (n=8, 99% confidence level)
Compound
Spiked
Concentration
(µg/L)
%
Recovery RSD (%)
MDL
(µg/L)
MCAA 2.0 105 4.1 0.5
MBAA 2.0 96 4.1 0.2
DCAA 2.0 97 4.1 0.2
TCAA 2.0 94 3.7 0.2
BCAA 2.0 94 3.2 0.2
BDCAA 2.0 94 3.3 0.2
DBAA 2.0 99 3.8 0.2
CDBAA 2.0 102 4.5 0.3
TBAA 2.0 81 15.3 0.7
y = 9.7142x - 1.5579 R² = 0.98
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0 10.0
Concentr
ations (
µg/L
)
Ratio of peak areas
BDCAA
y = 11.721x - 3.0322 R² = 0.9699
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0 10.0
Concentr
ations (
µg/L
)
Ratio of peak areas
CDBAA
y = 13.781x - 2.7168 R² = 0.9832
0
20
40
60
80
100
120
0.0 2.0 4.0 6.0 8.0
Concentr
ations (
µg/L
)
Ratio of peak areas
TBAA
182
Figure 10-5 GC-ECD method for MCAA- spike recovery chart
Figure 10-6 GC-ECD method for MBAA- spike recovery chart
2009 2010
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
MBAA
2009 2010
130%
70%
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
MCAA
2009 2010
130%
70%
183
Figure 10-7 GC-ECD method for DCAA- spike recovery chart
Figure 10-8 GC-ECD method for TCAA- spike recovery chart
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
DCAA
2009 2010
130%
70%
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
TCAA
2009 2010
130%
70%
184
Figure 10-9 GC-ECD method for DBAA- spike recovery chart
Figure 10-10 GC-ECD method for BCAA- spike recovery chart
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
DBAA
2009 2010
130%
70%
0
20
40
60
80
100
120
140
0 10 20 30 40
Sp
ike
re
cove
ry (
%)
Sample set
BCAA
2009 2010
130%
70%
185
Flame Atomic Absorption Spectrometry 10.2.3
Figure 10-11 Example of FAAS calibration curves for iron, copper and lead
y = 13.813x - 0.4067 R² = 0.9974
0
2
4
6
8
10
0 0.2 0.4 0.6 0.8
Me
tal C
on
ce
ntr
ation
(m
g/L
)
Absorbance
Fe
y = 5.8834x - 0.0437 R² = 0.998
0.0
0.5
1.0
1.5
2.0
2.5
0 0.1 0.2 0.3 0.4
Me
tal C
on
ce
ntr
ation
(m
g/L
)
Absorbance
Cu
y = 18.423x + 0.0167 R² = 0.9911
0.0
1.0
2.0
3.0
4.0
5.0
0.00 0.10 0.20 0.30
Me
tal C
on
ce
ntr
atio
n
(mg/L
)
Absorbance
Pb
186
Table 10-3 FAAS method – Method detection limits for iron, copper and lead
Compound Spiked
Concentration
(mg/L)
%
Recovery RSD (%)
MDL
(mg/L)
Iron 0.2 101 6.7 0.041
Copper 0.1 95 3.3 0.009
Lead 0.1 95 8.3 0.024
187
Ion Chromatography 10.2.4
Figure 10-12 Example of IC calibration curves for chloride, sulphate and bromide
y = 1E-06x + 1.2651 R² = 0.9995
0
20
40
60
80
100
120
0.0E+00 2.0E+07 4.0E+07 6.0E+07 8.0E+07 1.0E+08
Ion
Con
ce
ntr
ation
(m
g/L
)
Area counts
Chloride
y = 0.0035x + 2.6704 R² = 0.9969
0
40
80
120
160
0 10,000 20,000 30,000 40,000 50,000
Ion
Con
ce
ntr
ation
(µ
g/L
)
Area counts
Bromide
y = 2E-06x - 2.0752 R² = 0.9803
0
20
40
60
80
100
0.0E+00 1.0E+07 2.0E+07 3.0E+07 4.0E+07
Ion
Co
nce
ntr
atio
n (
mg/L
)
Area counts
Sulphate
188
Figure 10-13 Example of IC calibration curves for nitrite, nitrate and phosphate
y = 7.5E-07x - 1.2E-02 R² = 1.0E+00
0.0
0.2
0.4
0.6
0.8
1.0
1.2
0.0E+00 5.0E+05 1.0E+06 1.5E+06
Ion
co
nce
ntr
ation
(m
g/L
)
Area counts
Nitrite
y = 5.3E-07x + 1.4E-02 R² = 1.0E+00
0
2
4
6
8
10
12
0.0E+00 5.0E+06 1.0E+07 1.5E+07 2.0E+07
Ion
co
nce
ntr
atio
n (
mg/L
)
Area counts
Nitrate
y = 1.7E-06x + 1.2E-02 R² = 1.0E+00
0.0
0.2
0.4
0.6
0.8
1.0
1.2
0.0E+00 2.0E+05 4.0E+05 6.0E+05 8.0E+05
Ion
co
nce
ntr
ation
(m
g/L
)
Area counts
Phosphate
189
Table 10-4 IC method – Method detection limits for nitrite, nitrate, chloride, bromide, sulphate
and phosphate
Compound Spiked
Concentration
(mg/L)
%
Recovery RSD (%)
MDL
(mg/L)
Nitrite 0.1 95 4.1 0.012
Nitrate 1.0 101 1.9 0.057
Chloride 0.02 91 14 0.007
Bromide 0.02 99 4.2 0.003
Sulphate 1.0 90 10 0.28
Phosphate 0.05 131 10 0.02
References
Domino, M. M., Pepich, B.V., Munch, D.J., Fair, P.S. and Xie. Y. (2003) Method 552.3:
Determination of haloacetic acids and dalapon in drinking water by liquid-liquid
microextraction, derivatization, and gas chromatography with electron capture detection.
Revision 1.0. Technical Support Center, Office of Ground Water and Drinking Water,
U.S. Environmental Protection Agency. Cincinnati, Ohio 45268.
Munch, J. W., and Bassett, M. V. (2004) Method 521: Determination of nitrosamines in drinking
water by solid phase extraction and capillary column gas chromatography with large
volume injection and chemical ionization tandem mass spectrometry (MS/MS). Version
1.0. National Exposure Research Laboratory, Office of Research and Development, U.S.
Environmental Protection Agency, Cincinnati, Ohio 45268.
190
10.3 Free Chlorine 24-hour Residuals through the Duration of Metal
Coupon Conditioning
Figure 10-14 24 hour chlorine residuals monitored through the duration of metal coupon
conditioning
0
3
6
9
12
0 5 10 15 20HO
Cl 24h r
esid
ual (m
g/L
)
Time (d)
Initial dosage Fe Fe+ortho
0
3
6
9
12
0 5 10 15 20HO
Cl 24h r
esid
ual (m
g/L
)
Time (d)
Initial dosage Cu Cu+ortho
0
3
6
9
12
0 5 10 15 20HO
Cl 24h r
esid
ual (m
g/L
)
Time (d)
Initial dosage Pb Pb+ortho
191
10.4 HAA Speciation for Fresh Metal Coupons in Britannia Water
Figure 10-15 HAA speciation with time in the presence 1 mg/L orthophosphate for fresh metal
coupons with Britannia Water in one set of experiments, error bars indicate the measured
maximum and minimum values (n=2)
0
10
20
30
40
50
0 10 20 30
HA
A (
µg
/L)
Time (hours)
Fe
DCAA
TCAA
MCAA
0
10
20
30
40
50
0 20 40 60 80
HA
A (
µg
/L)
Time (hours)
Cu
0
10
20
30
40
50
0 20 40 60 80
HA
A (
µg
/L)
Time (hours)
Pb
0
10
20
30
40
50
0 20 40 60 80
HA
A (
µg
/L)
Time (hours)
Control
192
10.5 XRD Analysis for Fresh Iron Coupons
Figure 10-16 Comparison of XRD patterns for oxidized iron coupon in the presence of polyphosphate-
orthophosphate blends and polished iron coupon as control.
Po
ly-o
rth
o b
len
ds
10 20 30 40 50 60 70 80
Inte
nsity
Ca
rbo
n
Fe
Fe Fe C
on
tro
l
193
10.6 Metal Release Kinetics and Results for Metal Surface Analysis
Figure 10-17 Kinetics of metal release from fresh metal coupons in the presence and absence of
corrosion inhibitors with NH2Cl in Mannheim Water; disinfectant concentrations, 12.3 mg/L;
error bars indicate the measured maximum and minimum values (n=2)
0
1
2
3
0 20 40 60 80 100
Pb (
mg/L
)
Time (hours)
Pb
Pb+polyphosphate
Pb+orthophosphate
0
2
4
6
8
10
0 5 10 15 20 25 30
Fe (
mg/L
)
Time (hours)
NH2Cl
Fe
Fe+polyphosphate
Fe+orthophosphate
0
1
2
2
3
4
0 20 40 60 80 100
Cu (
mg/L
)
Time (hours)
Cu
Cu+polyphosphate
Cu+orthophosphate
194
Figure 10-18 Kinetics of metal release from fresh metal coupons with time in the presence and
absence of corrosion inhibitors in Britannia Water; disinfectant concentrations, 12.3 mg/L; error
bars indicate the measured maximum and minimum values (n=2)
0
4
8
12
16
0 10 20 30
Fe (
mg/L
)
Time (hours)
HOCl
Fe
Fe+poly
Fe+ortho
0
2
4
6
0 20 40 60 80
Cu(m
g/L
)
Time (hours)
Cu
Cu+poly
Cu+ortho
0
2
4
6
8
0 20 40 60 80
Pb (
mg/L
)
Time (hours)
Pb
Pb+poly
Pb+ortho
0
4
8
12
16
0 10 20 30
Fe (
mg/L
)
Time (hours)
NH2Cl
Fe
Fe+poly
Fe+ortho
0
2
3
5
6
0 20 40 60 80
Cu (
mg/L
)
Time (hours)
Cu
Cu+poly
Cu+ortho
0
2
4
6
8
0 20 40 60 80
Pb (
mg/L
)
Time (hours)
Pb
Pb+poly
Pb+ortho
195
Figure 10-19 Kinetics of metal release from corroded coupons in the absence/presence of
orthophosphate in Britannia Water, initial disinfectant concentrations 5.5 mg/L; error bars
indicate the measured maximum and minimum values (n=2)
0
2
4
6
0 10 20 30 40 50 60
Fe (
mg/L
)
Time (hours)
HOCl
Fe Fe+ortho
0.0
0.5
1.0
1.5
2.0
0 10 20 30 40 50 60
Cu (
mg/L
)
Time (hours)
Cu Cu+ortho
0
1
2
3
4
0 10 20 30 40 50 60
Pb (
mg/L
)
Time (hours)
Pb
Pb+ortho
0
2
4
6
0 10 20 30 40
Fe (
mg/L
)
Time (hours)
NH2Cl
Fe
Fe+ortho
0.0
0.5
1.0
1.5
2.0
0 10 20 30 40 50 60
Cu (
mg/L
)
Time (hours)
Cu Cu+ortho
0
1
2
3
4
0 10 20 30 40 50 60
Pb (
mg/L
)
Time (hours)
Pb Pb+ortho
196
Figure 10-20 Comparison of elemental distribution for iron, copper and lead coupons in the
absence and presence of orthophosphate with HOCl
7%
47%
6%
37%
3%
Fe+HOCl
Fe
O
Ca
C
Si 2p
15%
52%
3%
23%
5% 2%
Fe+ortho+HOCl
Fe
O
Ca
C
P
Si 2p
28%
0%
16%
53%
1% 2%
Cu+HOCl
C
Ca
Cu
O
Cl
S
11%
7%
10%
58%
11%
3%
Cu+ortho+HOCl
C
Ca
Cu
O
P
Cl
31%
22%
47%
Pb+HOCl
Pb
C
O
6%
13%
56%
6%
9%
10%
Pb+ortho+HOCl
Pb
C
O
Al
Ca
P
197
10.7 NDMA Formation in the Presence of Orthophosphate in Modified
Pipe Loops
Figure 10-21 NDMA formation in four pipe loops in the presence of orthophosphate under two
flow conditions
0
10
20
30
40
2 6 12
ND
MA
(ng/L
)
HRT (hours)
Turbulent
0
10
20
30
40
2 6 12
ND
MA
(ng/L
)
HRT (hours)
Laminar
198
10.8 Effects of Corrosion Inhibitors and the Extent of Metal Corrosion
on Monochloramine Degradation and NDMA Formation
This research investigated the impacts of corrosion inhibitors and metal corrosion on
monochlormaine decay in two water matrices. However, due to consistently low formation of
NDMA in these two water matrices, the effects of phosphate-based corrosion inhibitors on
NDMA formation kinetics could not be evaluated in the two tested water matrices (Waters A and
B) when in contact with iron, copper and lead. Although not being completed to the same extent
as Chapter 3 and other chapters, some preliminary results concerning NDMA formation from
pre-dosed DMA in Milli-Q water are provided in this section in order to give implications about
the impacts of corrosion inhibitors and metal materials on NDMA formation.
199
The objective of this section was to investigate the effects of pipe materials (ductile iron,
copper and lead) and phosphate-based corrosion inhibitors (orthophosphate and polyphosphate)
on NH2Cl degradation and NDMA formation by performing material-specific formation
potential (MS-FP) and material-specific simulated distribution system (MS-SDS) tests at bench
scale. It is hypothesized that the application of phosphate-based corrosion inhibitors which
significantly affect corrosion rates may, in turn, impact NH2Cl degradation and NDMA
formation.
Materials and Methods 10.8.1
10.8.1.1 Reagents and materials
All chemicals used in this study were ACS grade or higher. The chlorine dosing solution
(approximately 3500 mg/L as Cl2) was prepared by diluting a concentrated solution of sodium
hypochlorite (NaOCl, 6%, VWR) in Milli-Q water. The NH2Cl dosing solution was then
prepared by adding the chlorine dosing solution to a well-stirred ammonium chloride solution
(700 mg/L as N). The chlorine dosing solution and ammonium chloride solution were combined
at a Cl2/N molar ratio of 0.8:1 to achieve a NH2Cl solution. Then the NH2Cl solution was
equilibrated for at least 30 min before use. The concentration of NH2Cl was then measured by
the indophenol method with a Hach DR 2700 spectrophotometer, and it ranged between 1400
and 1600 mg/L. At least 90% of the added chlorine was converted to NH2Cl. The water in batch
reactors (one-liter capacity amber bottles with PTFE lined caps) was dosed with preformed
NH2Cl.
Due to the consistently low formation of NDMA in two selected water matrices and to
ensure NDMA formation at a measurable level, one NDMA precursor, dimethylamine (DMA),
was spiked into Milli-Q water at a concentration of 1 µg/L to evaluate the impacts of corrosion
inhibitors and metal corrosion on NDMA formation.
The dosing solution of orthophosphate (Na3PO4) or polyphosphate ([Na(PO2)]6) was
prepared at a concentration of 500 mg/L as P and kept in dark at 4ºC. The targeted dosages for
each of corrosion inhibitors in test solution was 1 mg/L as P. Test coupons of ductile iron, copper
and lead were purchased from Metal Samples Co., Alabama, US. The size of these coupons is
1 2”3”1 1 ”. Unchlorinated post-filtration water was collected from two water treatment
200
plants in Ontario (Mannheim Water and Britannia Water). Water quality parameters are listed in
Table 10-5.
Table 10-5 Water quality parameters for post-filtration water
Parameters Values
Mannheim Water Britannia Water
pH 7.5 ± 0.2 6.6±0.5
Alkalinity (mg/L) 187±33 12±2
UV254 (cm-1
) 0.058±0.008 0.065±0.005
TOC (mg/L) 4.5 ± 0.3 3.3±0.2
SUVA(cm-1
·L/mg) 0.013±0.002 0.019±0.0002
Bromide (μg L) 65.0±15.5 15.8±0.2
Chloride (mg/L) 84.5 ± 2.5 3.1±0.6
Sulfate (mg/L) 35.0 ± 2.0 23.8±0.8
Cl-:SO4
2- ratio 2.4 ± 0.1 0.13±0.04
10.8.1.2 Experimental procedures
Experiments were performed with both fresh and pre-corroded metal coupons. Fresh
metal coupons were prepared before each batch of experiments by removing any corrosion
products from the coupons by polishing with 60-grit sandpaper followed by 120-grit sandpaper,
and then rinsing with deionized water and acetone followed by Milli-Q water. These polished
coupons simulated new pipes without any impacts from service age. The corroded metal coupons
were prepared by soaking them in tap water in the absence and presence of orthophosphate for at
least two weeks to allow corrosion products to be built up on the surface (details are provided in
Section 3.2.1). Orthophosphate was selected since it is in common use and it can effectively
control metal release in a short term for fresh metal coupons, especially for copper and lead.
“Material-specific” simulated distribution system or MS-SDS tests incorporated the
influences of the distribution system pipe wall on disinfectant residual stability and DBP
generation. Details of MS-SDS procedures have been described by Brereton and Mavinic (2002).
They consist of incubating metal coupons in water samples under conditions representative of
actual field conditions in terms of reaction time, pH, temperature, and disinfectant application. In
experiments with corroded metal coupons, MS-SDS tests were performed and initial NH2Cl
concentration was approximately 5.5 mg/L. In experiments with fresh metal coupons, material-
specific formation potential (MS-FP) tests were performed similarly to MS-SDS tests, except
that MS-FP tests applied a higher concentration of NH2Cl (12.3 mg/L) than would be
201
encountered during typical water treatment. This was done to meet the high chlorine demand of
fresh metal coupons and their initial corrosion products and ensure detectable NH2Cl residuals
after 24 hours. All of the MS-SDS and MS-FP experiments were performed in 1 L amber bottles
(chlorine-demand free) with PTFE lined caps and at room temperature (21±2 ºC). No pH
adjustment was performed, and thus reactions of metal corrosion and NH2Cl degradation in the
two water matrices proceeded under their respective ambient conditions. All of the metal
coupons were suspended in amber bottles using nylon threads. To maximize the contact of water
with coupon surface, a Big Bill Orbital shaker was used (Barnstead International) to maintain a
gentle mixing at a speed of 25 rpm.
Two sets of kinetic experiments were performed to test the reproducibility of the results,
and duplicate tests were conducted in each set of experiments. All of the tests also included
control samples, which were prepared similarly to test samples but without metal coupons. To
account for possible differences in NDMA yields due to the different pH conditions of the water
that were tested, a small number of additional experiments was performed, in which the pH was
controlled at 8.3±0.2 using 1 mM borate buffer in these tests, and the initial NH2Cl concentration
was set 14.5±0.5 mg/L as Cl2. Analytical methods that were employed to examine water quality
and the analytes of interest are summarized in Table 10-6.
Table 10-6 Summary of analytical methods
Analyte Unit Instrument /procedure Reference method
TOC mg/L OI-Analytical TOC analyzer SM1 5310 C
pH pH meter
UVA254 cm-1
Hewlett Packard 8452A Diode
Array UV spectrophotometer
SM 5910B
Alkalinity mg/L Titration SM 2320B
NH2Cl mg/L Hach DR2700 Spectrophotometer Hach method 101712
Anions µg/L Dionex DX-300 Series Ion
Chromatography System
SM 4110 B
NDMA ng/L Varian 4000 GC-MS USEPA 5213
Metal (Fe, Cu and Pb) mg/L Varian SpectrAA.20 SM 3111 B
Notes: 1. SM represents Standard Methods for the Examination of Water and Wastewater (APHA, AWWA, WEF, 2005); 2.
Hach, 2007; 3. Munch and Bassett, 2004.
202
Monochloramine Degradation 10.8.2
10.8.2.1 Fresh Coupons in Two Water Matrices
In the two investigated water matrices, NH2Cl degradation generally followed first-order
kinetics. The overall decay rate constants of NH2Cl (k) were obtained by fitting a pseudo-first-
order decay equation to the measured NH2Cl residual concentrations from each experiment. k is
usually considered as the sum of a first-order bulk decay constant (kb) and a first-order wall
decay constant (kw) (Rossman et al., 1994). In the present tests, the bottles had been made
chlorine demand free prior to testing, so the walls of the bottles did not contribute significantly to
the overall chlorine demand and the bulk water reaction in the MS-PF tests could be considered
to be the same as would occur in control samples without metal coupons. Thus in these tests, the
wall decay constant was the difference between the overall decay constant and the bulk decay
constant obtained from the control samples. Table 10-7 summarizes NH2Cl overall decay
constants (k) and the kb values for the three tested metal materials in the absence and presence of
corrosion inhibitors. An overview of the results for Mannheim Water, for example, is provided in
Figure 10-22 to illustrate the effects of metal materials and corrosion inhibitors on NH2Cl
degradation in Mannheim Water (as determined by the difference in bar size relative to their
respective control samples) in addition to showing the general bulk water decay (represented by
the control samples themselves).
Table 10-7 Comparison of NH2Cl decay constants for fresh coupons in two water matrices (n=4)
Material Corrosion
inhibitors Mannheim Water
Britannia
Water
k
Iron
None 0.1298 ±0.0073 0.1460±0.0438
Orthophosphate 0.1100 ±0.0151 0.1802±0.0423
Polyphosphate 0.2092 ±0.1099 0.2844±0.0021
Copper
None 0.0150 ±0.0086 0.0509±0.0098
Orthophosphate 0.0093 ±0.0016 0.0144±0.0003
Polyphosphate 0.0131 ±0.0062 0.0259±0.0054
Lead
None 0.0075 ±0.0012 0.0110±0.0006
Orthophosphate 0.0066 ±0.0003 0.0097±0.0015
Polyphosphate 0.0094 ±0.0037 0.0122±0.0008
kb Water
None 0.0081 ±0.0008 0.0159±0.0018
Orthophosphate 0.0074 ±0.0008 0.0122±0.0001
Polyphosphate 0.0105 ±0.0032 0.0141±0.0003
Initial NH2Cl concentration: 12.3 mg/L as Cl2
203
Figure 10-22 NH2Cl overall decay constants for fresh coupons with Mannheim Water; NH2Cl
12.3 mg/L, error bars indicate the measured maximum and minimum values (n=2)
For each metal material, a single-factor ANOVA test at a confidence level of 95% was
applied to the overall decay constants to determine whether the treatment factor (corrosion
inhibitors) had a significant influence on the response factor (NH2Cl decay constants). When the
ANOVA test signified statistically significant impacts of corrosion inhibitors on NH2Cl decay, a
Fisher’s Least Significant Difference (LSD) test was further applied to determine if significant
differences existed between each pair of treatments at a 95% of confidence level (Montgomery,
2000).
As shown in Table 10-7 and Figure 10-22, kb for Mannheim Water was not significantly
affected by corrosion inhibitors. This is consistent with the previous finding that corrosion
inhibitors did not significantly impact disinfectant bulk water degradation (Zhang and Andrew,
2012). Small variations in kb values also show good QA/QC between experiments. The most
striking feature of the data shown in Figure 10-22 is that iron was more much reactive with
NH2Cl than copper and lead. NH2Cl decay constants in the presence of iron coupons were two
orders of magnitude higher than the bulk water reaction constants irrespective of the types of
corrosion inhibitors present. Therefore, pipe wall reactions dominated NH2Cl consumption in the
presence of iron coupons (i.e., reactions with the coupon surface or with iron released from the
coupons). Corrosion inhibitors did not exhibit significant impacts on NH2Cl decay for iron
coupons (p value >0.05, the ANOVA test).
0.00
0.05
0.10
0.15
No inhibitors Orthophosphate Polyphosphate
1st o
rde
r d
eca
y c
on
sta
nt (h
-1)
Fe Cu Pb Control
0.21±0.08
kb
kw k
204
The NH2Cl overall decay rates for fresh copper coupons were 11~ 80% higher than those
in the bulk water (Figure 10-22), indicating that the contributions of the wall reactions and the
bulk reactions to NH2Cl overall decay were on a similar scale. Results of the single factor
ANOVA tests demonstrate that the addition of phosphate-based corrosion inhibitors did not
significantly impact NH2Cl degradation in the presence of copper coupons (p value >0.05).
Interestingly, the overall NH2Cl decay constants for lead coupons were similar to those for the
bulk water irrespective of the types of corrosion inhibitors. As a result, wall decay constants of
lead coupons were calculated to be negligible. Corrosion inhibitors did not significantly affect
NH2Cl decay for lead coupons as well (p value >0.05).
In order to help differentiate between effects due to reactions with the coupons directly
and effects due to reactions with ions released from the coupons, dissolved metal concentrations
at 24 hours were measured and presented in Figure 10-23. The released metal concentrations
were much higher for the iron coupons than for the lead and copper coupons, suggesting that the
higher NH2Cl decay rate for test solutions containing iron coupons could be attributed to
reactions with iron that was released from the coupons. Comparable iron concentration at 24
hours in the absence and presence of corrosion inhibitors (Figure 10-23) indicates that
phosphate-based corrosion inhibitors had no significant impacts on iron corrosion in Mannheim
Water. The variability in the decay constants for iron coupons in the presence of polyphosphate
Figure 10-23 Released metal concentrations at 24 hours for fresh metal coupons in Mannheim
Water; NH2Cl 12.3 mg/L; error bars indicate the measured maximum and minimum values (n=2)
0
2
4
6
8
No inhibitors Orthophosphate Polyphosphate
Me
tal co
nce
ntr
ation
(m
g/L
)
Fe Cu Pb
205
was large (between 0.13 h-1
and 0.29 h-1
, Figure 10-22), likely due to different polishing
conditions of the iron coupon surface between batches of experiments.
For copper coupons, orthophosphate effectively decreased the released copper
concentration from 1.5 mg/L in the absence of corrosion inhibitors to 0.7 mg/L, and reduced the
released lead concentration from 0.7 mg/L in the absence of corrosion inhibitors to 0.2 mg/L. For
copper and lead coupons in contact with NH2Cl over short period oxidation, Cu2O and divalent
lead solids have been reported as primary corrosion products (Vasquez et al., 2006; Xiao et al.,
2007). As such, the decrease in copper and lead concentrations due to the addition of
orthophosphate can stoichiometrically reduce the consumption of NH2Cl by 0.9 mg/L for copper
and 0.3 mg/L for lead. These two values only account for 4% and 0.8% of the initial NH2Cl
concentration applied (12.3 mg/L) for copper and lead, respectively. As a result, the slight
decrease in NH2Cl consumption due to the reduced copper and lead corrosion rates by
orthophosphate would not be expected to result in a significant difference in NH2Cl decay
relative to that without the addition of orthophosphate. Therefore, it was concluded that released
copper and lead did not contribute significantly to NH2Cl decay in this water matrix.
The relative sequence of reactivity of the three metal materials with NH2Cl in Britannia
Water had a similar pattern to that observed in Mannheim Water (Table 10-7) except that
phosphate-based corrosion inhibitors significantly affected NH2Cl decay for fresh iron and
copper coupons (p values of 0.04 and 0.02, respectively). Results of the LSD tests indicate that
polyphosphate statistically increased NH2Cl decay for iron coupons, while both polyphosphate
and orthophosphate statistically decreased NH2Cl degradation for copper coupons. Figure 10-24
displays dissolved iron and copper concentrations at 24 hours for Britannia Water. For iron
coupons, released iron concentrations were significantly increased after the addition of
phosphate-based corrosion inhibitors likely due to the formation of iron-phosphate complexes
and thus the increased iron solubility (McNeill and Edwards, 2000). Increased NH2Cl
degradation with increases in released iron concentrations may also suggest that the reactions
with released iron from the coupons were a primary reason for the accelerated NH2Cl decay for
iron coupons in Britannia Water. For copper coupons, results of stoichoimetrical calculation
indicate that the significant decrease in NH2Cl decay after the addition of phosphate-based
corrosion inhibitors was primarily due to the effective reduction of released copper
concentrations.
206
Figure 10-24 Released iron and copper concentrations at 24 hours in the absence and presence of
corrosion inhibitors for fresh metal coupons in Britannia Water; NH2Cl 12.3 mg/L; error bars
indicate the measured maximum and minimum values (n=2)
In addition, as shown in Table 10-7, NH2Cl degradation constants in Britannia Water in
contact with copper and lead as well as in bulk phase were significantly higher than those in
Mannheim Water (p-values 0.003, 0.02 and 0.0009, respectively), but were only statistically
different for fresh iron coupons in two water matrices at a confidence level of 85%. Edwards and
Dudi, 2004) have attributed the relatively low confidence levels in statistical analysis for iron to
low number of samples collected, which may also be applied to this study. Since the pH and
alkalinity values in Britannia Water were lower than those in Mannheim Water (Table 10-5), the
metal coupons were more subject to corrosion in Britannia Water which in turn would induce
more NH2Cl to participate into reactions with metal coupons. The details about the impacts of
water quality on metal corrosion are provided in Chapter 4. The accelerated auto-decomposition
of NH2Cl in the bulk Britannia Water at low pH may also be responsible for the increased NH2Cl
degradation than in Mannheim Water.
10.8.2.2 Corroded Coupons in Two Water Matrices
For corroded metal coupons, NH2Cl degradation also followed first-order kinetics. Table
10-8 summarizes NH2Cl overall decay constants (k) and the kb values for the three tested metal
materials in the absence and presence of orthophosphate. For both water matrices, the reactivity
of three metal materials with NH2Cl followed the sequence of Fe>CuPb. Only for corroded iron
0
3
6
9
12
No inhibitors Orthophosphate Polyphosphate
Me
tal co
nce
ntr
ation
(m
g/L
)
Fe Cu
207
coupons, NH2Cl degradation was dominated by wall reactions, while wall decay constants of
corroded copper and lead coupons were calculated to be negligible. In addition, orthophosphate
significantly increased NH2Cl degradation for the corroded iron coupons (ANOVA tests, p-value
of 0.04) with reasons as explained in Section 10.7.2.1, but exhibited no statistical impacts for the
corroded copper and lead coupons (p-values >0.05).
Table 10-8 Comparison of NH2Cl overall decay constants for corroded coupons between
Mannheim Water and Britannia Water (n=4)
Material Corrosion
inhibitors
Mannheim
Water
Britannia
Water
k
Iron None 0.0495±0.0071 0.0733±0.0158
Orthophosphate 0.0781±0.0037 0.1447±0.0142
Copper None 0.0114±0.0013 0.0186±0.0001
Orthophosphate 0.0065±0.0040 0.0173±0.0011
Lead None 0.0084±0.0011 0.0168±0.0013
Orthophosphate 0.0057±0.0009 0.0156±0.0006
kb Water None 0.0083±0.0010 0.0182±0.0008
Since Britannia Water had lower pH and alkalinity than Mannheim Water (Table 10-5), it
was expected to be more corrosive to metal coupons than Mannheim Water. Therefore, more
NH2Cl should have participated in the metal corrosion reactions, leading to more rapid NH2Cl
degraded in Britannia Water than in Mannheim Water. As shown in Table 10-8, corroded iron
coupons in contact with Britannia Water had k values of at least 48% higher than with Mannheim
Water, and the wall decay constants (kw, the difference between k and kb) at least 34% higher.
Therefore, NH2Cl decay for corroded iron coupons exhibited the expected trend. However, for
copper and lead, although NH2Cl decay constants in Britannia Water were significantly higher
than those in Mannheim Water with all of the p-values less than 0.05, the consumption of NH2Cl
by coupon surfaces (indicated by kw values) was calculated to be negligible. As such, the
increased NH2Cl overall decay in Britannia Water for both copper and lead coupons was
primarily attributed to its significantly increased bulk phase reaction rates (p-value of 0.008),
mostly likely via auto-decomposition.
NDMA Formation from DMA for Fresh Coupons 10.8.3
In addition to NH2Cl degradation, this section also investigated interactive impacts of
metal corrosion and corrosion inhibitors on NDMA formation. As discussed in Section 10.7.2.2,
208
orthophosphate did not significantly affect NH2Cl degradation for corroded copper and lead
coupons. Since NH2Cl concentration is a rate-limiting factor of NDMA formation, it was
postulated that orthophosphate would not affect NDMA formation for the corroded metal
coupons as well. Therefore, only fresh metal coupons were employed in this test to evaluate the
interactive impacts of metal corrosion and corrosion inhibitors on NDMA formation. In addition,
since corrosion inhibitors did not significantly affect NH2Cl degradation in the bulk water, only
NDMA formation in Milli-Q water in the absence of metal coupons and corrosion inhibitors was
considered as a control for comparison purposes.
Figure 10-25 compares NDMA formation at 24 hours for each combination of metal
materials and corrosion inhibitors. NH2Cl residuals at 24 hours are also demonstrated in Figure
10-25. Compared with bulk water control, Milli-Q water in contact with iron coupons had
significantly higher yields of NDMA irrespective of the types of corrosion inhibitors. In
particular, iron coupons exposed to polyphosphate had a NDMA yield at least seven folds as
high as that in bulk water control. NH2Cl residual at 24 hours for iron coupons in the presence of
polyphosphate dramatically decreased to 0.2 mg/L, but there were still 3.1 mg/L and 3.6 mg/L
NH2Cl remaining at 24 hours for iron coupons alone and in the presence of orthophosphate,
respectively.
Figure 10-25 NDMA formation from DMA as a result interactions of metal coupons, corrosion
inhibitors and NH2Cl. Initial NH2Cl 14.5±0.5 mg/L, DMA 1 µM/L, pH 8.3, in Milli-Q, error
bars indicate standard deviation (n=3)
0
4
8
12
16
0
600
1200
1800
2400
3000
Polyphosphate No inhibitors Orthophosphate
NH
2C
l (m
g/L
)
ND
MA
(n
g/L
)
Fe-NDMA Cu-NDMA Pb-NDMAWater-NDMA Fe-NH2Cl Cu-NH2ClPb-NH2Cl Water-NH2Cl
209
NDMA formation at 24 hours for copper coupons alone, and copper coupons treated with
polyphosphate and orthophosphate also increased compared with that in bulk water control.
However, the magnitude of the increase was not as significant as that for iron coupons. The
addition of orthophosphate decreased NH2Cl degradation, and 10.1 mg/L NH2Cl remained after
24 hours. However, there were only 8.2 mg/L NH2Cl left for copper coupons alone and 8.5 mg/L
for copper coupons in the presence of polyphosphate after 24 hours.
Lead coupons had comparable 24-hour NH2Cl residual concentration and NDMA
formation with bulk water control regardless of the presence of corrosion inhibitors. As
discussed in Section 10.7.2.1, NH2Cl wall reactions had little contribution to overall NH2Cl
decay for fresh lead coupons. It follows that the consumption of NH2Cl due to metal surface
reaction may not significantly affect the residual concentration of NH2Cl in reaction with DMA
in aqueous solution. As such, it is expected to observe comparable NDMA formation in the
water containing lead coupons and bulk water control.
To statistically evaluate the interactive effects of metal corrosion and corrosion inhibitors
on NDMA formation, and to determine whether the treatment factor (corrosion inhibitors) had a
significant influence on the response factor (NDMA concentrations), a single-factor ANOVA
test at a confidence level of 95% was applied to the results that were obtained for each metal
material combined with different corrosion inhibitors. When the ANOVA test results signified
significant difference in NDMA formation between metal coupons and bulk water control, each
treatment was further compared with bulk water control by using the LSD tests for single
contrasts at a confidence level of 95% (Montgomery, 2000). Results of the single contrasts using
the LSD test are summarized in Table 10-9.
Table 10-9 Summary of the results for the single contrasts using the LSD test
Comparison Significantly different NDMA formation
No inhibitor Polyphosphate Orthophosphate
Milli-Q Fe Yes Yes Yes
Milli-Q Cu Yes Yes No
Note: a statistically significant at a confidence level of 95%
The ANOVA test results confirmed that NDMA formation for iron and copper coupons
in the absence and presence of corrosion inhibitors was significantly different from that in bulk
water control. The results of single contrasts using the LSD tests for iron coupons indicate that
210
iron coupons significantly increased NDMA formation, regardless of the presence of and/or
types of corrosion inhibitors. According to NH2Cl degradation studies in Section 10.7.2.1, wall
reactions primarily due to iron corrosion dominated NH2Cl decay, leading to less availability of
NH2Cl residual to react with the constituents in the bulk phase. Since NH2Cl concentration is a
rate-limiting factor of NDMA formation, it follows that lower concentrations of NDMA should
have formed in the water in contact with iron coupons than in the bulk water control. The
unexpected elevation in NDMA formation in the water containing iron coupons suggests that
iron and/or iron corrosion products may catalyze the transformation of DMA to NDMA. To
identify the possible roles of iron corrosion products in NDMA formation, NDMA concentration
was further plotted versus released total iron concentration. As shown in Figure 10-26, NDMA
yields increased with increases in released total iron concentrations. It indicates that iron
catalysis during NDMA formation was positively affected by released iron concentration. It can
also be observed in Figure 10-26 that corrosion inhibitors significantly affected NDMA
formation primarily by affecting released iron concentration.
Figure 10-26 Total iron concentrations and NDMA formation for iron coupons in the absence
and presence of corrosion inhibitors, error bars indicate standard deviation (n=3)
In terms of copper coupons, the LSD comparison results indicate that NDMA formation
significantly increased for copper coupons alone and treated with polyphosphate compared with
that in bulk water control. However, there was no statistical difference in NDMA yields between
copper coupons exposed to orthophosphate and bulk water control. It was also observed that
0
2
4
6
8
10
0
600
1200
1800
2400
3000
Polyphosphate No inhibitors Orthophosphate
To
tal iro
n (
mg
/L)
ND
MA
(n
g/L
)
NDMA Total iron
211
NDMA concentration positively increased with increases in released Cu(II) concentration
(Figure 10-27). Therefore, it was hypothesized that copper can catalyze NDMA formation with
the extent of catalysis being affected by copper concentration. Further experiments have been
performed to test this hypothesis (Chapter 6). Results of these experiments confirmed copper
catalysis during NDMA formation, and that copper catalysis nonlinearly increased with
increasing copper concentration. Although 0.5 mg/L copper was released from copper coupons
treated with orthophosphate, its catalytic effect on NDMA formation may not be significant
under current experimental conditions so that comparable NDMA yields were observed for
copper coupons in the presence of orthophosphate and bulk water control.
Figure 10-27 Cu(II) concentrations and NDMA formation for copper coupons in the absence and
presence of corrosion inhibitors, error bars indicate standard deviation (n=3)
This study provides important implications about factors affecting NH2Cl stability and
NDMA formation in distribution mains and household plumbing. Phosphate-based corrosion
inhibitors have been increasingly used for corrosion control. However, these corrosion inhibitors
may exhibit either beneficial or detrimental effects on NH2Cl stability and NDMA formation,
depending on metal types, age and water quality. Therefore, their impacts on NH2Cl decay and
NDMA formation due to the interactions with metal surfaces should be considered site-
specifically.
0.0
0.3
0.6
0.9
1.2
0
200
400
600
800
1000
Polyphosphate No inhibitors Orthophosphate
Cu
(II
) (m
g/L
)
ND
MA
(n
g/L
)
NDMA Cu(II)
212
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Moffat, E. (2006) Formation and decay of disinfection byproducts in the distribution
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Brereton, J. A., and Mavinic, D. S. (2002) Field and material-specific simulated distribution
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Charrois, J. W. A., Boyd, J. M., Froese, K. L., and Hrudey, S. E. (2007) Occurrence of N-
nitrosamines in Alberta public drinking-water distribution systems. Journal of
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Choi, J., Duirk, S. E., and Valentine, R. L. (2002) Mechanistic studies of N-
nitrosodimethylamine (NDMA) formation in chlorinated drinking water. Journal of
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Dartmann, J., Alex, T., Dorsch, T., Schevalje, E., and Johannsen, K. (2004) Influence of
decarbonisation and phosphate dosage on copper corrosion in drinking water systems.
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Demora, S. J., and Harrison, R. M. (1984) Lead in tap water - contamination and chemistry.
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Duirk, S.E., Gombert, B., Croue, J.P. and Valentine, R.L. (2005) Modeling NH2Cl loss in the
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Facey, R. M., and Smith, D. W. (1995) Soft, low-temperature water-distribution corrosion:
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Gerecke, A.C. and Sedlak, D.L. (2003) Precursors of N-nitrosodimethylamine in natural waters.
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McNeill, L. S., and Edwards, M. (2000) Phosphate inhibitors and red water in stagnant iron
pipes. Journal of Environmental Engineering-Asce, 126(12), 1096-1102.
Maddison, L. A., Gagnon, G. A., and Eisnor, J. D. (2001) Corrosion control strategies for the
Halifax regional distribution system. Canadian Journal of Civil Engineering, 28(2), 305-
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Montgomery, D. C. (2000) Design and Analysis of Experiments, 5th edition, John Wiley &
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Environmental Protection Agency, Cincinnati, Ohio 45268, 2004.
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216
10.9 Degradation Potential of Iron, Lead and Their Corrosion Products
on HAA9
This research investigated the degradation potential of iron and lead on HAA9 in Milli-Q
water under controlled experimental conditions. However, no further experiments have been
conducted to evaluate the impacts of water quality of the two tested water matrices (Waters A
and B) on HAA degradation by iron and copper. Although not being completed to the same
extent as other chapters, some results of this section which deal with iron support the discussions
in Section 3.3.3, Chapter 3. The results that deal with lead degradation potential also help
understand of the fate of HAA compounds in lead pipes.
217
The purpose of this section was to investigate the degradation potential of iron and lead
as well as their corrosion products on HAA9 under controlled experimental conditions. Since iron
is widely used in water mains and lead may still be present in old lead service lines, soldered
joints and brass plumbing fittings, the knowledge about the degradation potential of iron and lead
for HAA9 is important to understand the fate of these HAA compounds in water distribution
mains and domestic plumbing systems.
Materials and Methods 10.9.1
All chemicals used in this study were ACS grade or higher. Test coupons of ductile iron
and lead were purchased from Metal Samples Co., Alabama, US. The size of these coupons was
1 2”3”1 1 ”. For the tests investigating the degradation potential of iron and its corrosion
products on HAAs, iron coupons were conditioned in the absence and presence of different
corrosion inhibitors in tap water for 24 hours to simulate the scenario when pipes have been
corroded over a short period. Initial chlorine concentration for conditioning was 12.3 mg/L.
Corrosion inhibitors that were applied were orthophosphate 1.0 mg/L as P, polyphosphate 1.0
mg/L as P, and orthophosphate-polyphosphate blends (0.5/0.5 mg/L as P). Borate buffer
(H3BO3/NaOH) solutions were also prepared to control the solution pH at desired levels
(8.3±0.2). In the tests investigating the degradation potential of lead, fresh lead coupons were
used to simulate the contact of DBPs with metal surface without any impacts from service age.
All of the tested metal coupons were suspended in 1 L amber bottles using nylon threads. To
maximize the contact of water with coupon surface, a Big Bill Orbital shaker was used
(Barnstead International) to maintain a gentle mixing at a speed of 25 rpm.
The degradation potential of lead corrosion products, including PbO, Pb(OH)2(CO3)2, and
PbO2 (Sigma Aldrich), on HAAs were also investigated. Each solid corrosion product was
spiked into 1 L Milli-Q water in the form of powder at a concentration of 1 g/L, and then mixed
with Milli-Q water thoroughly by a stir bar. After each prescribed reaction time, 50 mL water
was taken from reaction bottles and filtered through 0.2 µm Nylaflo® Nylon membrane filter
paper (Pall Corporation) to separate powdered corrosion products from the aqueous solution.
All of the experiments were performed in duplicate along with control bottles, which
were prepared in a similar manner except that no coupons were soaked or no corrosion products
218
were added. To account for any influence from filter paper and filtration performance on the
concentration of HAA9, control samples for the experiments employing powdered lead corrosion
products were also filtered under similar conditions and its HAA9 concentrations after filtration
were reported as Control in the results section.
In addition, X-ray diffraction (XRD) analysis was applied to the surface of iron coupons
after 24-hour conditioning to characterize the compositions of iron corrosion products after a
short-term reaction and to determine their possible roles during the degradation of HAA9. A
Siemens D5000 Diffractometer System operating at 50 kV/35 mA was used to collect the
diffraction patterns. A high-power, line focus Cu-K-source was used combined with a solid
state Kevex detector for elimination of K-lines. The experimental data were collected on a step
scan mode (0.02° /1.5 sec) within the most informative range (2-theta degrees). The obtained
data were processed by various Diffrac Plus software packages including Eva 8.0 and
Topas v. 2.1. For all of the tested iron coupons, due to the formation of non-uniform corrosion
layers on the coupon surface as a result of short-term reactions, XRD analysis was carried out
on powders which were scratched from the iron coupon surface.
Nine species of haloacetic acid (HAA9) were analyzed according to USEPA Standard
Method 6251 B (APHA et al., 2005). The principle was based on liquid–liquid extraction of
HAA9 with methyl-tert-butyl-ether (MtBE) at an acidic pH followed by diazomethane
derivatization and gas chromatography with electron capture detection (GC/ECD) analysis. The
surrogate standard was 2,3,5,6-tetrafluorobenzoic acid. All of the extracted and derivatized
samples were stored at -10 °C or less, and analyzed within 21 days of extraction. The gas
chromatograph used for this analysis was a Hewlett Packard 5890 Series II Plus Gas
Chromatograph. The GC column was DB1701 (30 m х 0.2 mm х 0.2 μm). The column
temperature programs for HAAs measurement were: hold at 35 °C for 10 min; ramp to 65 °C at
2.5 °C/min; ramp to 85 °C at 10 °C/min; ramp to 205 °C at 20 °C/min and hold for 7 min.
HAA9 Degradation by Corroded Iron Coupons 10.9.2
The degradation potential of iron was investigated by employing pre-corroded iron
coupons, and the variation of HAA9 concentrations with reaction time is shown in
219
Figure 10-28 Compared with the control bottles where no loss of HAA9 was observed,
HAA9 were rapidly degraded in the water containing corroded iron coupons. The HAA9
degradation kinetics followed the sequence of decreasing order: poly/ortho blends,
polyphosphate orthophosphate, and no corrosion inhibitor.
Figure 10-28 Reduction of HAA9 by corroded iron coupons in the absence and presence of
different corrosion inhibitors. Milli-Q water, error bars indicate the measured maximum and
minimum values (n=2)
The degradation of each species of HAA9 except for MCAA in the presence of iron
coupons treated with poly/ortho blends is plotted in Figure 10-29. All of the eight HAA
compounds were readily degraded by corroded iron coupons. DCAA degraded relatively slowly
Figure 10-29 HAA degradation by corroded iron coupons treated with poly/ortho blends, Milli-Q
water, error bars indicate the measured maximum and minimum values (n=2)
0
100
200
300
400
500
0 20 40 60 80 100
HA
A9 (
µg
/L)
Time (hours) Fe Fe+polyFe+ortho Fe+poly/orthoControl
0
10
20
30
40
50
0 20 40 60 80 100
HA
A (
µg
/L)
Time (hours)
MBAA DCAA
TCAA BCAA
DBAA BDCAA
CDBAA TBAA
220
in the earlier 24 hours. For MCAA, as shown in Figure 10-30, its concentrations increased with
increasing reaction time in the initial 24 hours, and then followed a decreasing trend after 24
hours.
Figure 10-30 Reduction of MCAA by corroded iron coupons in the absence and presence of
different corrosion inhibitors. Milli-Q water, error bars indicate the measured maximum and
minimum values (n=2)
Hozalski (2001) has proposed sequential hydrogenolysis to be a primary degradation
mechanism for HAA compounds in contact with Fe0. To identify if Fe
0 was present on the
surface of the pre-conditioned iron coupons, XRD analysis was performed, and the results are
shown in Figure 10-31. A significant amount of Fe0 (represented by black peak at 45 degree) was
present on the corroded iron surface even after 24 hour oxidation. Due to the presence of Fe0 on
Figure 10-31 Comparison of XRD patterns for oxidized iron coupon in the presence of
poly/ortho-phosphate blends and polished iron coupon as control. Red: iron coupon as control;
Black: non-scratched iron coupon with poly/ortho-phosphate blends
0
20
40
60
80
100
0 20 40 60 80 100
MC
AA
(µ
g/L
)
Time (hours)
Fe Fe+polyFe+ortho Fe+poly/orthoControl
Inte
nsity
10 20 30 40 50 60 70 80
2-theta (degree)
Fe
Fe Fe Fe
+poly
/ort
ho
Fe
contr
ol
221
the surface of freshly oxidized coupons, therefore, it was hypothesized that sequential
hydrogenolysis was the reaction pathways for HAA degradation in the presence of these
corroded iron coupons.
To test this hypothesis, single species degradation experiments were performed regarding
MCAA, DCAA, TCAA, and BCAA due to their wide presence in the chlorinated water. The
results of the degradation experiments for four compounds are collectively illustrated in Figure
10-32. The degradation products, including HAA compounds and anions (e.g. chloride and
bromide) were also analyzed. However, the concentrations of chloride and bromide were below
method detection limits (MDLs) of these anions by ion chromatography, and thus not reported
herein.
Figure 10-32 HAA speciation degradation by corroded iron coupons treated with poly/ortho
blends, error bars indicate the measured maximum and minimum values (n=2)
As shown in Figure 10-32, all of the parent HAA compounds experienced degradation in
the presence of these corroded iron coupons over the reaction period, whereas their
0
15
30
45
60
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
MCAA MCAA Control
0
10
20
30
40
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
DCAA
MCAA
DCAA control
0
10
20
30
40
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
TCAA DCAA
MCAA Control
0
10
20
30
40
0 20 40 60 80 100 120
HA
A (
µg
/L)
Time (hours)
MCAA
BCAA
BCAA control
222
concentrations in the control samples remained constant. For example, TCAA was reduced
within 72 hours with DCAA and MCAA being produced. DCAA concentrations were
consistently below 2 µg/L, and MCAA concentrations were initially increased to 7.6 µg/L at 44
hours and then dropped to 2 µg/L after 120 hours. The plots of TCAA degradation with reaction
time illustrate the destruction of the parent compound, the production of the reaction
intermediates (e.g. DCAA), and the degradation of final products (e.g. MCAA). Therefore, the
results of single species experiments suggest that HAA degradation by corroded iron coupons,
most likely due to the presence of Fe0, also followed the mechanism of sequential
hydrogenolysis. The degradation of these HAA compounds can be illustrated as follows
(Hozalski, 2001):
TCAADCAA
DCAAMCAA
MCAACAA
BCAAMCAA
In particular, during BCAA degradation, MCAA rather than MBAA was produced, indicating
that bromide was favored to be lost rather than chloride in the process of hydrogenolysis.
Constant concentrations of MCAA, DCAA, TCAA and BCAA in the control samples over the
reaction time indicate that sequential hydrogenolysis by iron was the only mechanism
responsible for the loss of HAA.
Generally, the degradation of HAA compounds followed the pseudo-first-order decay
kinetics. The pseudo-first-order degradation constants for four HAA compounds were estimated
by fitting a first-order decay equation to the HAA concentration data in Microsoft Excel. The
degradation rates for MCAA, DCAA, TCAA and BCAA were 0.015, 0.031, 0.047, and 0.055 h-1
,
respectively. Therefore, the reactivity of four HAA compounds with corroded iron coupons
followed a sequence of BCAA >TCAA >DCAA>MCAA, which is in agreement with the results
reported by Zhang et al. (2004). In their study, at similar initial concentrations of four HAA
compounds, the degradation rates of MCAA, DCAA, TCAA and BCAA were 0.001, 0.008,
0.502 and 3.87 h-1
, respectively. Since the pseudo-first-order rate of DBP reduction by iron was
223
significantly affected by pH, the mineral surface and iron surface speciation (Chun et al., 2005),
the difference in the decay constants for four compounds between this study and those reported
in the literature is likely because the reaction pH, iron type and/or iron surface conditions were
different.
In addition, according to the mechanisms of hydrogenolysis, DCAA can be formed as an
intermediate product of TCAA and BDCAA degradation. Therefore, it may initially accumulate
in the solution, and thus exhibit slower degradation compared with other compounds (Figure
10-29). Although MBAA and DBAA were formed as intermediate products as well, their
degradation rate constants were reported at least one order of magnitude higher than that of
DCAA (Zhang et al., 2004). Therefore, these two species experienced faster degradation rather
than accumulating in the system (Figure 10-32). The variation of MCAA concentrations over the
reaction time, as shown in Figure 10-32, indicates that MCAA, as an intermediate product and
due to its slow degradation kinetics, could also accumulate in the solution in the initial 24 hours
(illustrated by the increase of its concentration before 24 hours). After the degradation of other
HAA compounds was completed, the degradation of MCAA dominated its formation reactions
and caused the decrease in MCAA concentrations after 24 hours.
HAA Degradation by Lead 10.9.3
Compared with iron, a standard reduction potential of which being 0.44V at 25°C (Fe
Fe2+
+2e), lead has a lower standard reduction potential of 0.13V at 25°C (Pb Pb2+
+2e).
Although smaller, the positive reduction potential of lead determines that it may degrade HAA
compounds as well. Therefore, this study also investigated the degradation potential of lead for
HAA compounds, and the results are shown in Figure 10-33. Except for MCAA, the
concentrations of other species exhibited a decreasing trend with increasing reaction time.
DCAA had a relatively smaller degradation rate compared with other species, especially in the
initial 15 hours. In contrast, MCAA concentration increased quickly in the initial 10 hours (from
110 µg/L to 231 µg/L), and then tended to plateau afterwards (only 13% increase of MCAA
observed in the following 40 hours).
Single species degradation experiments by fresh lead coupons were also performed for
MCAA, DCAA, TCAA and BCAA. Figure 10-34 displays the variation of MCAA
224
concentrations with reaction time in the presence of fresh lead coupons. Constant concentrations
of MCAA over the test period of 48 hours indicate that lead may not reduce MCAA. While the
degradation of MCAA by iron was observed in Section 10.8.1, the difference in MCAA
degradation between iron and lead was primarily because lead has a weaker reductive activity
than iron.
Figure 10-33 Reduction of HAA9 by fresh lead coupons, Milli-Q water, pH 8.3, error bars
indicate the measured maximum and minimum values (n=2)
Figure 10-34 Degradation of MCAA by fresh lead coupons, Milli-Q water, pH 8.3, error bars
indicate the measured maximum and minimum values (n=2)
The degradation of DCAA, TCAA and BCAA in water samples containing fresh lead
coupons is illustrated in Figure 10-35. The pathways of DCAA, TCAA and BCAA degradation
were similar as those observed for corroded iron coupons. In the case of TCAA, its concentration
decreased over the reaction time with concomitant formation of DCAA and MCAA as
degradation products. DCAA concentration increased in the initial 10 hours, and then exhibited
0
60
120
180
240
300
0 10 20 30 40 50
HA
A (
µg
/L)
Time (hours)
TCAA BCAAMBAA MCAADBAA DCAABDCAA CDBAATBAA
0
30
60
90
120
150
0 0 1 2 2 4 8 10 24 33 48
MC
AA
(µ
g/L
)
Time (hours)
225
only 25% reduction in the following 22 hours. MCAA concentrations continuously increased
with increasing reaction time, and plateaued at 32 hours. The decreased concentration of TCAA
and the increased formation of DCAA and MCAA suggest that the degradation of HAA
compounds by lead also followed the mechanism of sequential hydrogenolysis. For BCAA, only
MCAA was detected as a degradation product rather than MBAA. It indicates that bromide was
preferentially removed from HAA compounds over chloride, which is similar to the observations
for corroded iron coupons. Constant concentrations of DCAA, TCAA and BCAA over the
Figure 10-35 Degradation of single HAA species by fresh lead coupons, Milli-Q water, pH 8.3,
error bars indicate the measured maximum and minimum values (n=2)
0
20
40
60
80
100
120
0 10 20 30 40
HA
A (
µg
/L)
DCAA MCAA
0
20
40
60
80
100
0 10 20 30 40
HA
A (
µg
/L)
TCAA
DCAA
MCAA
0
20
40
60
80
100
120
0 10 20 30 40
HA
A (
µg
/L)
Time (hours)
BCAA MCAA
226
reaction period in the control bottles (plots are not shown) indicate that hydrogenolysis was the
primary mechanism for the destruction of the parent HAA compounds, the production of the
reaction intermediates, and the accumulation of final products in the presence of lead.
The pseudo-first-order degradation rate constants for DCAA, TCAA and BCAA were
estimated by fitting a first-order decay equation to the HAA concentration data in Microsoft
Excel, and they were 0.048, 0.089, and 0.14 h-1
, respectively. As such, the reactivity of three
HAA species with lead followed a sequence in decreasing order of BCAA, TCAA, and DCAA.
The degradation of BCAA is relatively rapid likely because bromide was preferentially removed
relative to chloride by the tested lead coupons. However, due to different metal surface
conditions of iron and lead coupons tested in this study, direct comparisons of their degradation
rates for HAA compounds cannot be made.
The degradation potential of lead corrosion products, including PbO, Pb(OH)2(CO3)2 and
PbO2, on HAA9 were also investigated, and the results regarding MCAA, MBAA, DCAA,
TCAA and BCAA are shown in Figure 10-36. The concentrations of these compounds remained
constant over 72 hours. Similar trends were also observed for the other four HAA compounds. It
indicates that lead corrosion products, in which lead is in an oxidized state, did not have
degradation potential for HAA9, and only elemental lead (Pb0) could react with HAA
compounds.
Results of this study provide an important implication about the fate of HAA compounds
in distribution systems and household plumbing. Besides biodegradation, HAA compounds may
also experience abiotic degradation by metal materials, e.g., iron and lead. However, the loss of
HAA compounds due to the reduction by these pipe materials may only be viable in a pipe
environment where new pipe are installed and are not corroded substantially.
227
Figure 10-36 Degradation of HAA species in the presence of 1g/L lead corrosion products. Milli-
Q water, pH 8.3, error bars indicate the measured maximum and minimum values (n=2)
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0
30
60
90
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150
0 3 9 24 48 72
HA
A (
µg
/L)
Time (hours)
PbO
MCAA MBAA DCAA TCAA BCAA
0
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0 1 4 9 23 47 72
HA
A (
ug
/L)
Time (hours)
Pb(OH)2(CO3)2
MCAA MBAA DCAA TCAA BCAA
0
30
60
90
120
150
0 1 4 9 23 47 72
HA
A (
ug
/L)
Time (hours)
PbO2
MCAA MBAA DCAA TCAA BCAA
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