Article pubs.acs.org/est
Elevated Pb(II) Release from the Reduction of Pb(IV) Corrosion Product PbO2 Induced by Bromide-Catalyzed Monochloramine Decomposition Yuanyuan Zhang† and Yi-Pin Lin*,‡ †
Department of Civil and Environmental Engineering, Faculty of Engineering, National University of Singapore, Singapore 117576 Graduate Institute of Environmental Engineering, National Taiwan University, No. 1, Sec. 4, Roosevelt Road, Taipei 10617, Taiwan
‡
S Supporting Information *
ABSTRACT: The stability of Pb(IV) corrosion product PbO2 has been linked to lead contamination in chloraminated drinking water. Recent studies have shown that autodecomposition of monochloramine (NH2Cl) can cause lead release from PbO2 via reductive dissolution. Bromide (Br−) is a known catalyst for NH2Cl decomposition. In this study, we investigated whether Br−-catalyzed NH2Cl decomposition could further enhance lead release from PbO2. Our results showed that Br−_catalyzed NH2Cl decomposition did accelerate the reduction of PbO2, and the rate was enhanced by the lower pH value, higher Br−, and NH2Cl concentrations. A single linear correlation was found between the amount of NH2Cl decomposed and the amount of total Pb(II) released either in the absence or presence of Br−, suggesting that Br−catalyzed NH2Cl decomposition and NH2Cl autodecomposition may generate the same intermediate toward PbO2 reduction. The kinetics of total Pb(II) release can be successfully modeled by considering the overall rate of NH2Cl decomposition with NOH as the reactive intermediate responsible for PbO2 reduction. Our findings suggested that special attentions on lead contamination should be paid to systems with PbO2 scales and high Br−-containing source waters when switching disinfectant from free chlorine to monochloramine.
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via a series of reactions (reactions 1−14 in Table 1).27−29 Lin and Valentine17,20 investigated the reduction of PbO2 in the presence of NH2Cl and found that the amount of Pb(II) released from PbO2 was proportional to the amount of NH2Cl decomposed. It was proposed that an unknown intermediate formed from the autodecomposition of NH2Cl is capable of reducing PbO2 to cause Pb(II) release.17 We hypothesized that catalyzed NH2Cl decomposition may further enhance PbO2 reduction and, in turn, cause elevated Pb(II) release. Br− is a known catalyst for NH2Cl decomposition.30−33 In fresh water, Br− concentration usually ranges from trace amount to up to 0.5 mg/L.34 In seawater, Br− concentration is around 65 mg/L.35 Brackish water possesses a Br− concentration between the two extremes. After reverse osmosis treatment for drinking purpose, Br− concentration in desalinated seawater can be reduced to about 0.6 mg/L.36,37 Reactions governing Br−-catalyzed NH2Cl decomposition published in the literature are also shown in Table 1 (reactions 15−19 and 22). The rate of Br−-catalyzed NH2Cl decomposition is determined by the oxidation of Br− by NH3Cl+ (reaction 16).33 Thus, the rate of Br−-catalyzed
INTRODUCTION PbO2 has been identified in many distribution systems with a history of using free chlorine for disinfection.1−6 It is formed from the chlorination of lead-containing plumbing materials (LCPMs) with Pb(II) carbonate solids as the primary precursors.3,7−13 Recent studies have shown that PbO2 plays a critical role in regulating lead contamination in drinking water due to its insoluble nature and high oxidation potential.8 When PbO2 is formed and stable such as in the presence of free chlorine (HOCl/OCl−), it can serve as a sink for soluble lead. When PbO2 is detached or subject to reductive dissolution, it becomes a source of lead in drinking water. Changes in water chemistry that can alter the stability of PbO 2 may trigger high levels of lead released from PbO23,8,14−25 and cause serious lead contamination in drinking water. A notable case is the hazardous levels of lead found in Washington DC drinking water during the period of 2001− 2003 after the system switched disinfectant from free chlorine to monochloramine to control the formation of disinfection byproducts.3,13 This incident was believed to be triggered by the lower oxidation potential introduced by NH2Cl that destabilized PbO2 scales that had been formed on the lead pipe inner surfaces during the period when free chlorine was employed.3,13 NH2Cl is a weaker oxidant than free chlorine.26 It can autodecompose to form ammonia, chloride, and nitrogen © 2013 American Chemical Society
Received: Revised: Accepted: Published: 10931
October 29, 2012 August 26, 2013 August 28, 2013 August 29, 2013 dx.doi.org/10.1021/es402733e | Environ. Sci. Technol. 2013, 47, 10931−10938
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Table 1. Model for NH2Cl Autodecomposition, Br−-Catalyzed NH2Cl Decomposition and PbO2 Reduction no.
rate constant at 25 °C
reaction
reference
NH2Cl autodecomposition 1 HOCl + NH3 → NH 2Cl + H 2O
k1 = 1.5 × 1010 M−1 h−1
44
2
NH 2Cl + H 2O → HOCl + NH3
k2 = 6.8 × 10−2 h−1a
44
3
HOCl + NH 2Cl → NHCl 2 + H 2O
k3 = 1.0 × 10 M
4
NHCl 2 + H 2O → HOCl + NH 2Cl
k4 = 2.3 × 10−3 h−1
45
5
NH 2Cl + NH 2Cl → NHCl 2 + NH3
kdb
32
6
NHCl 2 + NH3 → NH 2Cl + NH 2Cl
k6 = 2.2 × 108 M−2 h−1
6
+
k7 = 4.0 × 10 M 5
−
7
NHCl 2 + H 2O → NOH + 2H + 2Cl
8
−
NOH + NHCl 2 → HOCl + N2 + Cl + H
10 11
forward reaction: HOCl ⎯→ ⎯ H + OCl
kfastc
backward reaction: H+ + OCl− → HOCl
kb = kfast/K
NH4 ↔ NH3 + H
+
−
+
kfast
backward reaction: NH3 + H+ → NH4 +
kb = kfast/K
H 2CO3 ↔ HCO3− + H+
backward reaction: 14
↔
+H
−1
−1
−1
−1
h
h
HCO3−
kfast
+
forward reaction:
backward reaction:
CO3−2
→
+
+H →
+H
39, 48 39, 48
49
49
49
c
pKa = 10.3
CO3−2
39, 47
kb = kfast/K
+ H → H 2CO3
+
HCO3−
46
39, 48
pKa = 6.3
forward reaction: H 2CO3 → HCO3− + H+
CO3−2
h
45
c
forward reaction: NH4 ⎯→ ⎯ NH3 + H+
HCO3−
−1
pKa = 9.3 + fast
13
h
pKa = 7.5
HOCl ↔ H + OCl
fast
12
k10 = 55.0 M
−
+
7
k9 = 3.0 × 10 M
NH 2Cl + NHCl 2 → N2 + 3Cl− + 3H+ +
−1
−1
k8 = 1.0 × 108 M−1 h−1
+
NOH + NH 2Cl → N2 + Cl− + H+ + H 2O
9
−1
kfast
+
49
c
kb = kfast/K
HCO3−
−
Br -catalyzed NH2Cl decomposition 15 NH 2Cl + H+ ↔ NH3Cl+
K = 28 M−1 kf = 2.16 × 10 M
forward reaction: NH 2Cl + H+ → NH3Cl+
6
backward reaction: NH3Cl+ → NH 2Cl + H+ 16
50 8
−1
h
−1
31
−1
kb = 7.71× 10 h
kBr = 1.6 × 108 M−1 h−1a
NH3Cl+ + Br − → NH3Br + + Cl−
c
30
17
NH 2Cl + NH3Br + ⎯→ ⎯ NHBrCl + NH4 +
kfast
18
HOCl + Br − → HOBr + Cl−
kHOCl = 5.1 × 105 M−1 h−1
51
kfastc
32
k20 = 7.2 × 105 M−1 h−1
this study
k21 = 5.0 × 108 M−1 h−1
this study
19 20 21 22
fast
fast
HOBr + NH 2Cl ⎯→ ⎯ NHBrCl + H 2O
NHBrCl + H 2O → NOH + 2H+ + Br − + Cl− −
NOH + NHBrCl → HOBr + N2 + Cl + H
+
kfast
fast
NHBrCl + NH 2Cl ⎯→ ⎯ N2 + Br − + 2Cl− + 3H+
proposed reactions for the formation of Pb(II) and reactions of secondary intermediate (NO2−) 23 NOH + PbO2 → Pb2 + + NO2− + OH− 24
−
NO2 + NH 2Cl + H 2O →
NO3−
+
25
NO2 + NHCl 2 + 2H 2O → HOCl +
NO3−
26
NO−2
NO3−
+ NHBrCl + 2H 2O → HOBr +
c
32
k23 = 1.3 × 105 m−2 h−1 k24 = 4.0 × 10 M
−1
+
−
k25 = 2.0 × 10 M
−1
+
−
k26 = 9.0 × 108 M−1 h−1
7
−
+ NH3 + H + Cl
−
32
+ NH3 + H + Cl
+ NH3 + H + Cl
8
−1
h
−1
h
this study this study this study this study
a Note: k2 and kBr (reactions 2 and 16) were found to significantly affect the rates of NH2Cl autodecomposition and Br−-catalyzed NH2Cl decomposition. Their values were calibrated in this study using the data obtained from control experiments. The determined k2 and kBr values were comparable to those reported by Morris and Isaac44 (7.6 × 10−2 h−1) and Trofe et al.30 (1.8 × 108 M−1 h−1), respectively. The values of k20, k21, k23, k24, k25, and k26 were determined by the least-square regression of experiment data. On the basis of sensitivity analyses, the values of k24, k25, and k26 can only be considered as approximate estimates. bkd = kH+[H+] + kH2CO3 + kHCO3− [HCO3−], where kH+ = 2.5 × 107 M−2 h−1, kH2CO3 = 4 × 104 M−2 h−1, and kHCO3− = 800 M−2 h−1. ckfast represents the rate constant of reaction proceeding at a very fast rate. A value of 1 × 1010 M−1 h−1 is used in the model calculations. This value ensures immediate equilibrium in weak acid−base reactions.52
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To account for the water-induced PbO2 reduction, control experiments using samples containing only PbO2 and water were conducted. In addition, experiments using solutions containing Br− but no NH2Cl were also performed as Br− was reported to be able to reduce PbO2 in acidic conditions.22 Preliminary NH2Cl-free experiments showed that Br− did not induce PbO2 reduction in neutral and alkaline pH values, but significant total Pb(II) released was observed at pH 6.0 when compared to water-only conditions (Figure S2, Supporting Information). This additional Pb(II) released should be taken into consideration when PbO2 reduction caused by Br−-catalyzed NH2Cl decomposition was investigated at this pH value. Duplicate experiments were conducted in selected conditions. The modeling of NH2Cl decomposition and total Pb(II) formation was conducted using MATLAB (Version 7.11.0.584 (R2010b)). Details of the model simulations are presented in the Supporting Information. Analytical Methods. Total Pb(II) concentration was measured using anodic stripping voltammetry (Metrohm VA 797 computrace) with 0.1 M acetate (pH 4.0) as the background electrolyte. The standard addition technique was employed. As the solution was acidified to pH 4.0, this method could capture all Pb(II) species including soluble Pb(II) ions and Pb(II) carbonate solids that may potentially precipitate in the experimental solutions.38 The presence of PbO2 did not cause interference in this method.17,20 NH2Cl concentration was determined using the iodometric method.17 The formed I3− concentration was quantified using a UV spectrometer (Shimadzu UV-1800) with a molar absorption coefficient of 23 325 M−1 cm−1 at 351 nm.17 AJEOL 6700F field emission scanning electron microscopy was used to acquire the SEM image of the synthesized PbO2, and a Siemens XRD D5005 was used to identify its mineralogy. The seven-point N2-BET specific surface area was determined by a NOVA 4200e surface area analyzer. The precalibrated pH meter (F-51, Horiba) was used to measure the solution pH value.
monochloramine decomposition is faster at a lower pH value, higher Br− concentration, and higher initial NH2Cl concentration.33 The objective of this study was to investigate the relationship between Br −-catalyzed NH2Cl decomposition and PbO2 reduction to examine our hypothesis, which is important for lead control in systems with high Br−-containing source waters when switching from free chlorine to NH2Cl for disinfection. Experiments were designed to determine whether the factors accelerating the rate of Br−-catalyzed NH2Cl decomposition can synchronically accelerate the rate of total Pb(II) release from PbO2. The relationship between the amount of NH2Cl decomposed and the amount of total Pb(II) released in the presence and absence of Br− was also determined to compare the reactivity of the intermediate formed from different decomposition pathways toward the reduction of PbO2. Finally, the identity of the intermediate was proposed, and the modeling of NH2Cl decomposition and total Pb(II) release from PbO2 in this system was attempted.
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MATERIALS AND METHODS Chemicals. Reagent grade chemicals were used in this study. PbO2 was synthesized by the chlorination of 1 mM Pb(NO3)2 (Sigma-Aldrich) with 1.1 mM free chlorine following the procedures described by Lin and Valentine.17 Scanning electron microscopy image (SEM) showed that synthesized PbO2 particles were aggregates consisting of nanosized particles (∼20−30 nm), and X-ray diffraction pattern (XRD) confirmed that this PbO2 was plattnerite (β-PbO2) (Figure S1, Supporting Information). Its specific surface area was determined to be 17.0 m2/g. NH2Cl solution was prepared by the addition of free chlorine to bicarbonate-buffered ammonium chloride solution (SigmaAldrich) with a Cl/N molar ratio of 0.7 according to published procedures.33 Free chlorine solution was prepared by diluting a NaOCl stock solution (∼4%, Sigma-Aldrich). Potassium bromide (Fisher Scientific) was used as the source of bromide. Sodium bicarbonate (Nacalai Tesque) was used as the source of dissolved inorganic carbon (DIC). One normal HCl and NaOH were used to adjust solution pH value. Ultrapure water obtained from a Millipore DirectQ system was used to prepare all experimental solutions. Lead Release Experiments. Experiments were conducted using 4 mM DIC buffered solutions with various concentrations of Br− (0.1−2.0 mg/L or 1.3−25.0 μM), NH2Cl (0.9−3.4 mg/ L as Cl2 or 12.9−48.3 μM), and pH value (pH 6.0 to 8.0) to investigate their impacts on total Pb(II) release from PbO2. The PbO2 loading employed was 10 mg/L. All experiments were conducted using 70 mL gastight polyethylene vessels at 25 °C. After being filled with the experimental solution and PbO2, the vessels were sealed without head space and covered by aluminum foil to prevent the decay of NH2Cl induced by light. These vessels were placed on an orbital shaker rotating at 200 rpm during the course of the experiment. One vessel was used each time for the determination of total Pb(II) and residual NH2Cl concentrations for a total period of 92 h. The total Pb(II) concentration (soluble + particulate) was measured using unfiltered sample to capture Pb(II) ions as well as Pb(II) carbonate solids that could potentially precipitate in the experiment.17 In selected experiments, soluble Pb(II) was measured using the samples passing through a syringe filter equipped with a 0.2 μm pore size PTFE membrane (Titan). In these measurements, particulate Pb(II) was determined as the difference between total Pb(II) and soluble Pb(II). The residual NH2Cl concentration was determined using the filtered samples.
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RESULTS AND DISCUSSION Synergistic Influences of Br− and NH2Cl on the Reduction of PbO2. The general influences of NH2Cl, Br−, and PbO 2 on the decomposition of NH 2 Cl and the corresponding release of total Pb(II) from PbO2 are shown in Figure 1. The addition of PbO2 slightly accelerated the rate of NH2Cl decomposition, and the addition of Br− accelerated the decomposition rate to a greater extent due to its catalytic effects. The greatest rate of NH2Cl decomposition was observed in the presence of both PbO2 and Br− (Figure 1a). In the absence of NH2Cl, Br− did not accelerate the reduction of PbO2 if compared to the control experiment without Br− at pH 7.0. NH2Cl, on the other hand, promoted the release of total Pb(II) from PbO2 (Figure 1b). These findings were consistent with previous studies showing that the reduction of PbO2 by Br− did not proceed appreciably at neutral to alkaline pH values22 and that the decomposition of NH2Cl could lead to the reduction of PbO2.17 When Br− and NH2Cl were present simultaneously, a higher rate of total Pb(II) release was observed. For example, after 92 h, total Pb(II) released from PbO2 in the presence of 6.3 μM Br− and 26.8 μM NH2Cl was 4.16 μM. When present individually, identical concentrations of Br− and NH2Cl only caused 0.59 and 2.56 μM of total Pb(II) released, respectively. Pb(II) ions released from PbO2 reduction may form Pb(II) carbonate solids considering the DIC level present in the experimental solutions. It was found that soluble Pb(II) contributed 44% and 27% to the total Pb(II) at pH 6.0 10933
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Figure 1. Synergistic influences of Br− and NH2Cl on (a) NH2Cl decomposition and (b) total Pb(II) release from PbO2 as a function of time. Experimental condition: 10 mg/L PbO2, DIC = 4 mM, pH 7.0, temperature = 25 °C.
Figure 2. Influence of pH value on (a) NH2Cl decomposition and (b) total Pb(II) release as a function of time. Experimental condition: PbO2 = 10 mg/L, initial NH2Cl concentration = 28.0 μM, Br− = 6.3 μM, DIC = 4 mM, temperature = 25 °C. Solid lines represent modeling results.
and pH 7.0, respectively. SEM images for lead solids phases collected at 48 and 92 h for pH 6.0 and pH 7.0 are shown in Figure S3 (Supporting Information). Before the experiment, aggregations of nanosized PbO2 solid were found (Figure S3a, Supporting Information)). At pH 6.0, plat-shaped hydrocerussite (Pb3(CO3)2(OH)2) was observed to coexist with PbO2 at 48 h (Figure S3b, Supporting Information) and bigger hydrocerussite with small rod-shaped cerussite (PbCO3) was found at 92 h (Figure S3c, Supporting Information). At pH 7.0, discrete hydrocerussite was found among PbO2 particles at 48 h (Figure S3d, Supporting Information), and more hydrocerussite but no cerussite was found at 92 h (Figure S3e, Supporting Information). It should be noted that only total Pb(II) can be used to fully account for the reduction of PbO2. Thus, total Pb(II) was used in the subsequent sections for detailed discussion. Influence of pH Value. The decomposition of NH2Cl and release of total Pb(II) from PbO2 as a function of time in Br−containing NH2Cl solutions at pH 6.0−8.0 are shown in Figure 2. The rates of both NH2Cl decomposition (Figure 2a) and total Pb(II) release (Figure 2b) increased with the decreasing pH value. The enhanced Pb(II) released from PbO2 at the lower pH value can be attributed to the accelerated Br−-catalyzed NH2Cl decomposition. The rate of Br−-catalyzed NH2Cl decomposition is controlled by the formation of NH3Cl+ and its subsequent oxidation by Br−.30−33 The formation of NH3Cl+ was enhanced at the lower pH value and resulted in the faster decomposition of NH2Cl as shown in Figure 2a. The concurrent dependence of Br−catalyzed NH2Cl decomposition and total Pb(II) released from
PbO2 on pH value was consistent with our hypothesis that the Br−-catalyzed NH2Cl decomposition can generate a higher concentration of the reactive intermediate that is capable of reducing PbO2. In the absence of NH2Cl, Br− itself with a concentration of 6.3 μM did not cause significant PbO2 reduction at neutral to alkaline pH values. At pH 6.0, the same concentration of Br− induced additional 0.40−0.72 μM of Pb(II) release at different reaction periods when compared to Br−-free conditions (Figure S2, Supporting Information). These additional Pb(II) release, however, only accounted for about 10% of total Pb(II) release induced by the Br−-catalyzed NH2Cl decomposition (Figure 2b), indicating that Br−-catalyzed NH2Cl decomposition significantly promoted PbO2 reduction at this slightly acidic pH value. Influence of Br− Concentration. The decomposition of NH2Cl and the release of total Pb(II) from PbO2 as a function of time in NH2Cl solutions containing 1.3−25.0 μM of Br− are shown in Figure 3. Results from the control experiment without Br− are also shown for comparison purposes. It was found that the rates of NH2Cl decomposition (Figure 3a) and total Pb(II) released from PbO2 (Figure 3b) increased with the increasing Br− concentration. The kinetics of Br−-catalyzed NH2Cl decomposition has been previously reported to follow a firstorder dependence on Br− concentration,32 which was also observed in our data. Our observation suggested that Br− indirectly enhanced the release of Pb(II) from PbO2 via Br−catalyzed NH2Cl decomposition. Influence of Initial NH2Cl Concentration. The decomposition of NH2Cl and the release of total Pb(II) from PbO2 as a function of time in Br−-containing NH2Cl solutions at pH 7.0 10934
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Figure 4. Influence of initial NH2Cl concentration on (a) NH2Cl decomposition and (b) total Pb(II) release as a function of time. Experimental condition: PbO2 = 10 mg/L, Br− = 6.3 μM, DIC = 4 mM, pH 7.0, temperature = 25 °C. Solid lines represent modeling results.
Figure 3. Influence of Br− concentration on (a) NH2Cl decomposition and (b) total Pb(II) release as a function of time. Experimental condition: PbO2 = 10 mg/L, initial NH2Cl concentration = 28.0 μM, DIC = 4 mM, pH 7.0, temperature = 25 °C. Solid lines represent modeling results.
with different initial NH2Cl concentrations are shown in Figure 4. The rates of NH2Cl decomposition (Figure 4a) and total Pb(II) released from PbO2 (Figure 4b) increased with the increasing initial NH2Cl concentration. In the control experiment without NH2Cl, 6.3 μM Br− caused 0.59 μM total Pb(II) released from PbO2 after 92 h, which was only 23% of the total Pb(II) released in the presence of 12.9 μM NH2Cl, the lowest NH2Cl concentration used in our experiments. Our results indicated that the release of Pb(II) from PbO2 was primarily attributed to NH2Cl decomposition, in which the kinetics was dependent on the initial NH2Cl concentration and could be significantly promoted by Br−. Relationship between NH2Cl Decomposition and Total Pb(II) Release. The relationship between the amount of NH2Cl decomposed and the amount of total Pb(II) released from PbO2 is presented in Figure 5, where all experimental data including those from NH2Cl autodecomposition (Figures S4 and S5, Supporting Information) and those from Br−-catalyzed NH2Cl decomposition (Figures 2−4) are shown. Linear relationships were observed for both Br−-free and Br−containing conditions suggesting that the release of Pb(II) from PbO2 is controlled by the decomposition of NH2Cl, most likely via the reactive intermediate formed during the decomposition. The slopes of the linear regressions with a 95% confidence level in the presence and absence of Br− were 0.237(±0.030) and 0.255(±0.021), respectively. There was no statistically significant difference between the two slopes,
Figure 5. Summary plot of NH2Cl decomposition vs total Pb(II) formation in all experimental conditions.
suggesting that the reactive intermediate generated from NH2Cl autodecomposition and Br−-catalyzed NH2Cl decomposition could be identical. The slope obtained from the regression of all data was 0.241(±0.021), suggesting that the 10935
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decomposition of 1 μM of NH2Cl would cause approximately 0.241 μM of total Pb(II) release from 10 mg/L of the PbO2 used in this study. Modeling of NH2Cl Decomposition and Total Pb(II) Release. The release of total Pb(II) from PbO2 in NH2Cl solutions either in the absence or presence of Br− could be attributed to the reaction between PbO2 and the reactive intermediate formed from NH2Cl decomposition. We assume that both NH2Cl decomposition pathways produced the same intermediate that can reduce PbO2 based on our observation shown in Figure 5 and the discussion presented above. In NH2Cl autodecomposition, NOH is an intermediate formed from the reaction between NHCl2 and H2O considering the mass and electron balances of the reaction.39 NOH can further react with NHCl2 and NH2Cl to accelerate NH2Cl decomposition (reactions 7, 8, and 9 in Table 1). In the presence of Br−, as an analogy to reaction 7, NOH could be formed from the reaction between NHBrCl and H2O (reaction 20 in Table 1) considering the similar redox property of NHCl2 and NHBrCl. NOH can be redox reactive because of the “+1” oxidation state of N. Thus, the reactive intermediate that can reduce PbO2 is proposed to be NOH. It should be noted that the reduction of PbO2 by NOH is a surface reaction which may involve several steps, including the adsorption of NOH on PbO2 surfaces, electron transfer between NOH and surface Pb(IV), possible formation of Pb(III) intermediate and NOH radical, and the release of final redox products including Pb2+ and NO2−. It is believed that the reactions of surface electron transfer are extremely fast and the concentrations of species involved are extremely low under steady state.22,24,40 Considering the whole reaction scheme shown in Table 1, these reaction steps are not likely to be rate-determining to affect the model simulation. Thus, an overall reaction was used to represent the formation of Pb(II) from the reduction of PbO2 by NOH (reaction 23 in Table 1). Conceptually, consumption of NOH due to PbO2 reduction could decrease the rate of NH2Cl decomposition. However, our results showed that the rate of NH2Cl decomposition in PbO2containing solution were slightly higher than those without PbO2 (Figure 1), suggesting that a secondary intermediate could be formed from PbO2 reduction and it could also react with NH2Cl, NHCl2, and NHBrCl, respectively. This secondary intermediate is proposed to be NO2−. In the reduction of PbO2 by NOH, two electrons are transferred from NOH to PbO2 to form NO2− and Pb2+. NO2− has been reported as the product formed from the oxidation of NOH.41 It is proposed that NO2− can be further oxidized by NH2Cl, NHCl2, and NHBrCl to form NO3− as shown in reactions 24, 25, and 26 in Table 1. The oxidation of NO2− by NH2Cl to form NO3− (reaction 24) has been previously investigated.33 Since reduction of PbO2 is a surface reaction, its rate should be proportional to the available surface area.16,22 Thus, the rate of total Pb(II) released from PbO2 (reaction 23 in Table1) can be expressed as the following: r = k 23(SA)[NOH]
model. The rate constants of reaction 2 and 16, which were found to be sensitive to NH2Cl decomposition, were calibrated using data collected from control experiments without PbO2. The dependences of rate constants and equilibrium constants on temperature were also considered when available.32 In total, we used 204 data points from eight sets of experiments (Figures 2−4 and Supporting Information Figures S4, S5, and S7) to determine six unknown rate constants. The determined rate constants are shown in Table 1, and the modeling of NH2Cl decomposition and total Pb(II) formation are shown in Figures 2−4. The apparent rate constant (k23) was determined to be 1.3 × 105 m−2 h−1. Reactions involving NHBrCl (reactions 20, 21, and 22) proceed faster than their counterparts involving NHCl2 (reaction 7, 8, and 10) due to Br substitution. Faster reaction rates due to Br substitution were also observed for the reactions of HOBr and HOCl when reacting with NH2Cl (reactions 3 and 19) and in disinfection byproducts formation.42,43 In general, the model provided fairly good simulations of experimental data, except that at pH 6.0 the NH2Cl decomposition was overestimated and total Pb(II) released was underestimated. These discrepancies may be attributed to the facts that NH2Cl decomposition model was established under neutral to alkaline conditions and may not be directly applicable to pH 6.032 and that the release of total Pb(II) resulting from Br−-induced PbO2 reduction at this pH value (Figure S2, Supporting Information) was not taken into consideration in the modeling. The concentrations of the two proposed intermediates, NOH and NO2−, obtained from model simulations under various conditions are shown in Figures S8−S11 in the Supporting Information. The concentration of NOH ranged from 6.0 × 10−6 to 4.0 × 10−3 μM and that of NO2− ranged from 1.2 × 10−5 to 9.0 × 10−4 μM. Model results also showed that the concentration of NOH was higher at a lower pH value and it did not accumulate over time, suggesting that the stagnation time of chloraminated water in contact with PbO2 would be the important factor leading to total Pb(II) release rather than the total water age. To examine the sensitivity of the proposed model to rate constants of reactions 23−26, we modeled NH2Cl decomposition and total Pb(II) release as a function of time by either increasing or decreasing these rate constants. The results showed that when k23 (reaction 23) was increased or decreased by 1 order of magnitude from 1.3 × 105 to 1.3 × 106 or 1.3 × 104, no significant change was found in NH2Cl decomposition; while the amount of total Pb(II) release approximately increased by 30% and decreased by 70%, respectively (Figure S12 in the Supporting Information). There were almost no changes found in NH2Cl decomposition and total Pb(II) release when the rate constants of reactions 24−26 were either increased or decreased by a factor of 100 (Figure S13 in the Supporting Information). The sensitivity analysis suggested that the primary cause for NH2Cl decomposition does not involve PbO2, presumably due to the low concentration of the formed intermediates and that total Pb(II) release is governed by reaction 23. Since the model outputs were not sensitive to reactions 24−26, their rate constants provided in Table 1 can only be considered as approximate estimates. Reactions 24−26 could lead to the formation of NO3−, which can be measured for more accurate determination of these rate constants in future studies. Environmental Implications. NH2Cl can alter the stability of PbO2 via the intermediate formed from its autodecomposition. In this study, Br−-catalyzed NH2Cl decomposition was
(1) −1
where r denotes the rate of PbO2 reduction (M h ), k23 denotes the apparent rate constant (m−2 h−1). SA denotes the available surface area of PbO2 (m2). MATLAB (Version 7.11.0.584 (R2010b)) was used to determine the unknown rate constants (reactions 20, 21, and 23−26) via the least-squares regression of experimental data and to simulate the concentrations of all species involved in the 10936
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demonstrated to be able to accelerate this process causing elevated total Pb(II) release. The intermediate responsible for PbO2 reduction was proposed to be NOH. Theoretically, the catalytic effect of Br− can exist as long as Br− is present. The significance would depend on water parameters including pH value and NH2Cl and Br− concentrations. As a bench mark, at pH 7.0 and NH2Cl = 2 mg/L as Cl2, the catalytic effect can be observed at a bromide concentration as low as 0.1 mg/L. NOM in drinking water is also known to be able to reduce PbO2 and cause additional Pb(II) release from PbO2.14,15,19,20 The amount of Pb(II) released is dependent on the pH value, NOM characteristics/concentration and the degree of NOM oxidation by disinfectant. It was reported previously that 2.5 mg/L (as C) of Iowa River NOM can cause 4.21 μM total Pb(II) release at pH 7.0, DIC = 1 mM, and PbO2 = 10 mg/L after 7 days.20 For the same PbO2 loading, pH value, and a shorter reaction period (92 h), 4.16 μM total Pb(II) release was observed under 28 μM NH2Cl and 6.3 μM Br− (Figure 2). These two conditions showed comparable total Pb(II) released. It should be noted that, however, Pb(II) release caused by the NH2Cl decomposition and NOM are not additive as NH2Cl can oxidize NOM to consume some of its reductive moieties that can reduce PbO2.20 For distribution systems with historically installed leadcontaining plumbing materials, particularly those in arid and coastal regions with higher bromide concentrations in their source waters, precaution on lead contamination should be taken if considering switching disinfectant from free chlorine to NH2Cl. Systems using reverse osmosis desalinated seawater, which may contain a Br− concentration up to 0.6 mg/L,37 may also pose potential risk of lead release, although the risk could be reduced by lowing the Br− concentration after mixing with conventionally treated water in the distribution system. Our results also indicated that stagnation time of chloraminated water in contact with PbO2, instead of total water age, governs total lead release from PbO2. This highlights the importance of implementing lead control strategies to prevent lead release from lead service lines and premise plumbing into stagnant water.
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ASSOCIATED CONTENT
S Supporting Information *
Additional information including model simulations, the SEM and XRD pattern of synthesized PbO2, SEM images of lead solids phase before and after experiments, the NH2Cl decomposition and the Pb(II) release as a function of time in control experimental conditions, simulated NOH and NO2− concentrations under different pH values as well as k23, k24, k25, and k26 influences on the NH2Cl decomposition and the Pb(II) release. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Phone: 886-2-3366-9314; fax: 886-2-2392-8830; e-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The authors thank financial support from National University of Singapore (R-302-000-049-112). 10937
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