Environ. Sci. Technol. 2009, 43, 60–65
UV Photodegradation of Inorganic Chloramines JING LI AND ERNEST R. BLATCHLEY III* School of Civil Engineering, Purdue University, West Lafayette, Indiana 47907-2051
Received June 13, 2008. Revised manuscript received October 28, 2008. Accepted October 28, 2008.
The ultraviolet (UV) photolysis of monochloramine (NH2Cl), dichloramine (NHCl2), and trichloramine (NCl3) in aqueous solution was investigated at wavelengths of 222, 254, and 282 nm. All three chloramines can be degraded by UV irradiation, and the quantum yields for these processes are wavelengthdependent. Stable photoproducts include nitrite, nitrate, nitrous oxide, and ammonium. Solution pH was observed to have little effect on the rate of photodecay; however, the product distribution showed strong pH dependence. Nitrate formation was favored at low pH, while nitrite formation was favored at high pH. The effects of pH on formation of N2O and NH4+ were less clear. On the basis of the results, a mechanism of photodecay of monochloramine is proposed.
Introduction Chlorine is the most commonly used disinfectant for potable water and recreational water. However, important drawbacks of chlorine-based treatment applications have been identified, including generation of disinfection byproduct (DBPs), some of which may express toxicity toward humans. Among chlorinated DBPs, inorganic chloramines (monochloramine (NH2Cl), dichloramine (NHCl2), and trichloramine (NCl3)) are common in waters that contain reduced nitrogen in the form of ammonia/ammonium. Moreover, inorganic chloramines are intentionally produced in some drinking water applications by the process of chloramination. It is well-known that inorganic chloramines are produced by substitution reactions between free chlorine and ammonia (see reactions 1-3). The reactions compete with each other and are dependent on the initial chlorine to nitrogen (Cl:N) ratio and pH. NH3 + HOCl T NH2Cl + H2O
(1)
NH2Cl + HOCl T NHCl2 + H2O
(2)
NHCl2 + HOCl T NCl3 + H2O
(3)
These reactions reach equilibrium very rapidly. In the pH range of 6.5-8.5, monochloramine is predominantly formed at chlorine to ammonia molar ratio e1. The so-called “breakpoint” for reactions between free chlorine and ammonia-nitrogen occurs at a molar ratio of approximately 1.6-1.7; the breakpoint corresponds to the Cl:N ratio at which ammonia-nitrogen is (almost) completely oxidized by free chlorine to nitrogen gas, nitrite, nitrate and other products. At Cl:N ratios beyond the breakpoint, residual chlorine will exist predominantly in the form of free chlorine, with smaller quantities of NCl3. * Corresponding author phone: (765)494-0316; fax: (765)496-3145; e-mail:
[email protected]. 60
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Monochloramine is a relatively stable oxidant and a weak disinfectant. It has often been used as a secondary disinfectant when free chlorine residuals were difficult to maintain in a distribution system (1). For chloramination, utilities generally apply a Cl:N molar ratio between 0.6 and 1 (2) to balance the risk of trihalomethane/haloacetic acid formation. Monochloramine is the predominant (+1-valent) residual chlorine species under typical drinking water chloramination conditions. Ammonia is not the only precursor to produce inorganic chloramines. Previous work has demonstrated that trichloramine is a common byproduct of chlorination of aqueous solutions containing organic-nitrogen compounds that are found in recreational water settings, including urea, creatinine, L-histidine, and L-arginine (3). Analysis of swimming pool water samples has demonstrated that trichloramine is common to most swimming pools, with a representative aqueous-phase concentration of approximately 0.1 mg/L (as Cl2). In the United States, free chlorine concentrations of 1-3 mg/L (as Cl2) (1.4-4.2 × 10-5 M) are recommended to ensure disinfection in swimming pools. Previous studies suggest that the NH4+ concentration in swimming pools is typically on the order of 10 µg/L (5.6 × 10-7 M) (4). Therefore, most swimming pools operate at a chlorine:ammonia-nitrogen molar ratio that is well beyond the breakpoint (with respect to NH4+-N), and as a result, trichloramine is formed. Trichloramine has been reported to function as an irritant to the eyes and upper respiratory tract, and it may contribute to acute lung injury in accidental, occupational, or recreational exposures to chlorine-based disinfectants (5, 6) in swimming pools. UV irradiation has emerged as an important alternative to chlorination, in part because it has been found to be very effective for inactivation of the (oo)cysts of Cryptosporidium and Giardia, two pathogenic microorganisms of major importance for the safety of drinking water (7) that are resistant to free chlorine and chloramines. At doses and wavelengths used in common disinfection applications, UVbased processes have also been demonstrated to be costcompetitive with chlorine and to generate little or no DBPs. Moreover, UV irradiation has been demonstrated to be effective for photolysis of some chemical constituents, thereby allowing for the possibility of improving water quality in terms of both microbiological and chemical composition. Among the chemicals that have been observed to be susceptible to UV photodegradation are the inorganic chloramines. This has potentially important implications in water treatment operations where UV and chlorine are used in combination. For example, many drinking water utilities plan to use UV irradiation and chloramination as complementary disinfection processes, with UV being used as a primary disinfectant, and chloramination being implemented to yield a stable disinfectant residual. In swimming pool settings, UV-based systems are being used with increasing regularity for purposes of improving disinfection and for control of chloramines. In fact, UV systems are being sold for chloramine control in swimming pool settings, despite the fact that the kinetics and mechanisms of the reactions that lead to chloramine decay have not been well-defined. Given this situation, existing designs of UV systems that have been implemented for chloramine control have been based on empiricism, rather than fundamental principles of photochemical reactor design. As such, a need exists to improve our understanding of the basic chemistry and 10.1021/es8016304 CCC: $40.75
2009 American Chemical Society
Published on Web 11/26/2008
photochemistry of UV-induced photodecay of inorganic chloramines. The objective of this research was to explore the photochemical behavior of inorganic chloramines. UV irradiation was accomplished with several different collimated UV sources, including low-pressure Hg (254 nm); a KrCl excimer lamp (222 nm); and a XeBr excimer lamp (282 nm). These sources allowed for characterization of wavelength dependence in the resulting photochemical reactions. The ability of UV irradiation to degrade these compounds was investigated. Membrane introduction mass spectrometry (MIMS) was employed to monitor the decay of chloramine compounds when subjected to germicidal UV radiation. An important benefit of analysis by MIMS is that it does not suffer from analytical interference to the same degree as conventional wet-chemistry methods used for monitoring of chemical disinfectants (e.g., titration or colorimetry) and DBPs (e.g., gas chromatography). MIMS measurements were accompanied by ion chromatography (IC) to provide a more detailed representation of the products of these photoreactions. On the basis of product identification by MIMS and IC, as well as previously published information regarding related reactions, a reaction mechanism is hypothesized to describe the photodecay of monochloramine. It is believed that dichloramine and trichloramine will decay by similar pathways.
Experimental Section Materials and Methods. All chemicals used in this study, unless otherwise noted, were reagent-grade, purchased from Sigma-Aldrich, and used without further purification. Dilution to target aqueous-phase concentrations was accomplished with distilled, deionized water. Free chlorine stock solutions and standard solutions of inorganic chloramines were prepared and measured in the same manner as described previously (8). 15NH2Cl was prepared by addition of NaOCl to a solution of 15NH4Cl using analogous conditions. The analysis of standard solutions of inorganic chloramines prepared in the manner described above showed that NH2Cl can be produced as an aqueous solution that contains negligible quantities of NHCl2 and NCl3; NHCl2 standards always contained small amounts of NH2Cl and NCl3; NCl3 can be produced with minimal quantities of NH2Cl and NHCl2, but free chlorine will always be present in solution with NCl3. N2O gas was prepared by the boiling water bath method (9). MIMS was employed to develop the calibration curve by monitoring the peak at m/z 44 for 14N2O. For experiments that were conducted with 15N-labeled precursors, the concentration of 15N2O was determined by abundance at m/z 46. It was assumed that 14N2O and 15N2O yielded similar mass spectra, with the exception of the mass shift of 2 amu for the different N isotopes. The use of 15N in photodecay experiments was necessary to avoid possible interference in the MIMS analysis by CO2. Apparatus. The MIMS system was based on a modification of an HP 5892 benchtop GC/MS comprising an HP 5972A mass selective detector (MSD) equipped with electron (70 eV) ionization (EI). Mass spectrum scan mode (28 e m/z e 200) coupled with EI was used to identify possible DBPs, while selected ion monitoring (SIM) mode was used for quantification. Other details of the configuration and setup for MIMS system and operational conditions can be obtained from ref 8. The concentrations of inorganic chloramines were determined by comparison of ion abundance measurements with those developed from a series of standard solutions. Ions at m/z 53, 87, and 119 amu were monitored for quantification of NH2Cl, NHCl2, and NCl3, respectively. Ions at m/z 30 and 46 amu were monitored for quantification of
15
N2 and 15N2O, respectively. All compound identifications by MIMS were confirmed by analysis of standard solutions. A Dionex 500DX ion chromatograph (IC), equipped with an AS19 anion separatory column and an AG19 guard column was employed to quantify the anion products formed in the process of photodecay. A second IC system was configured with a CD 25 conductivity detector, GP50 gradient pump, CS 17 column, and CG17 guard column to quantify ammonium, and sodium sulfite solution was added to the samples before injection to reduce monochloramine to ammonium (10). Experimental Procedure. Photolysis experiments were carried out in stoppered UV-transparent (quartz) cuvettes to prevent losses of chloramines and reaction products by volatilization, while allowing exposure to UV radiation. UV exposures were conducted using collimating devices to deliver UV radiation. The collimating devices allowed delivery of a uniform, nearly parallel (collimated) beam of radiation to the solution surface. The UV sources included a lowpressure Hg lamp, a KrCl excimer lamp, and a XeBr excimer lamp; these lamps each produced essentially monochromatic collimated beams, with characteristic wavelengths of 254, 222, and 282 nm, respectively. A calibrated radiometer (International Light Technologies, Peabody, MA) was employed to measure the incident radiation intensity. UV dose delivery to the sample was controlled by exposure time. Dose quantification accounted for reflection, absorbance, and beam divergence. Experiments generally involved exposure to UV doses up to approximately 400 mJ/cm2, corresponding to a maximum exposure time of 40-100 min, depending on the collimated source. Additional details of the experimental setup and the methods of data interpretation are provided in the Supporting Information (see Figure S1 and related text).
Results and Discussion Absorption Spectra of Chloramines. The molar absorptivities of NH2Cl, NHCl2, and NCl3 in aqueous solution are illustrated in Figure 1. The λmax of monochloramine was 245 nm, with a molar absorptivity 461 M-1cm-1; absorbance was weaker at the wavelengths involved in this study. Dichloramine and trichloramine are relatively strong absorbers at 222 nm; both compounds absorb radiation less efficiently at wavelengths of 254 and 282 nm. The molar absorptivities of chloramine compounds did not change substantially over the pH range 6.5-8.5 that was used in this study. The fact that inorganic chloramines absorb germicidally active UV radiation implies that it is possible for these compounds to undergo direct photochemical reactions as a result of exposure to radiation at wavelengths of 222, 254, and 282 nm (UV222, UV254, and UV282, respectively). Chloramine Photolysis. The photodecay of mono-, di-, and trichloramine was studied as a function of UV dose at pH 7.5 with UV222, UV254, and UV282 (Figure 2). After each exposure period, a sample was collected and the concentration of the target compound was measured by MIMS. The results indicated that all three inorganic chloramines are susceptible to photodecay by exposure to germicidal UV radiation, with strong wavelength dependence. Less than 5% of monochloramine was observed to be degraded after exposure to UV282 at dose ) 390 mJ/cm2, whereas roughly 30% of monochloramine photodecayed as a result of exposure to UV222 and UV254 at similar doses. In general di- and trichloramine were more sensitive to UV irradiation than monochloramine, and both compounds showed stronger wavelength-dependence in their reactivity. For example dichloramine showed roughly 99% decay as a result of exposure to a UV222 dose of approximately 343 mJ/cm2, whereas exposure of dichloramine to UV254 and UV282 doses of similar magnitude yielded decays of roughly 35% and 60%, respectively. Similarly, trichloramine showed decay of VOL. 43, NO. 1, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. Molar absorptivity of mono-, di-, and trichloramine in aqueous solution. Inset table lists molar absorptivities (M-1cm-1) of target compounds for the wavelengths investigated in this work. roughly 97% at a UV222 dose of 366 mJ/cm2, whereas trichloramine exposure to similar UV254 and UV282 doses yielded decays of approximately 70% and 65%, respectively. Values of the (apparent) quantum yield for photodecay of target compound (φλ) were estimated by nonlinear fitting of eq 4 to the data sets: ln
I0,λ [T] L ) - (1 - Raq)(1 - Rqw) ε · ln(10)φλ · t (4) [T]0 Eλ L+l λ
( )
(see Supporting Information for development of this equation and definition of terms). Fits of eq 4 to the measured values of [T] as a function of UV dose are included as solid lines in Figure 2. The model fit the data quite well, with R2 values greater than 0.98 in most cases (see Table 1). The lowest value of R2 was observed in the data set that defined photodecay of NH2Cl by UV282. The low R2 value in this case was attributed to the fact that there was very little loss of NH2Cl for this reaction condition. For all combinations of target compound and wavelength, experiments were repeated two or three times. In each case, the estimates of quantum yield based on these experiments varied by less than 5% (data not shown). In the dark, mono- and trichloramine are relatively stable at pH 6.5-8.5, such that changes in the concentration of these compounds due to “dark” reactions that occurred over the period of UV exposure (less than 100 min) could be ignored (data not shown). On the other hand, dichloramine is unstable at pH > 7.5. It can disproportionate to form monochloramine and trichloramine, or undergo autodecomposition (11). The decay of dichloramine in the dark is illustrated in Figure 2 for comparison with the rate of photodecay. Apparent quantum yields for UV-induced photodecay of inorganic chloramines are summarized in Table 1. For six of the nine conditions examined in this research, the quantum yield was greater than 1.0. This implies that UV-induced photodecay of inorganic chloramines may involve chain reactions. The quantum yield for photodecay of monochloramine has been reported previously by other researchers. Cooper et al. (12) reported a quantum yield of 0.26 mol/ Es for this reaction at wavelengths of 253.7 and 280.4 nm. Similarly, Watts and Linden (13) reported a quantum yield for this reaction of 0.3 mol/Es at 254 nm. In many respects, the experiments that were used to define these earlier 62
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FIGURE 2. UV photodecay of monochloramine (top), dichloramine (middle), and trichloramine (bottom) as the function of UV dose for three wavelengths of exposure at pH 7.5. The secondary horizontal axis refers to the dark decay experiment with NHCl2. Solid lines represent nonlinear fits of eq 4 to the data.
TABLE 1. Quantum Yield (mol/Es) for Photodegradation of Inorganic Chloramines under UV Irradiation (Values in Parentheses Represent the Standard Error of the Quantum Yield Estimate, and the R2 Value for Each Regression, Respectively) compound
282 nm
254 nm
222 nm
NH2Cl NHCl2 NCl3
0.21 ((0.068, 0.625) 2.25 ((0.086, 0.995) 9.50 ((0.350, 0.993)
0.62 ((0.049, 0.983) 1.80 ((0.073, 0.995) 1.85 ((0.114, 0.995)
3.46 ((0.228, 0.938) 2.26 ((0.128, 0.997) 0.58 ((0.039, 0.996)
quantum yield estimates were similar to the experiments reported herein; however, some notable differences were evident. For example, the previously reported investigations were based on aqueous solutions that included large stoichiometric excess of ammonia-nitrogen relative to residual chlorine. In contrast, the experiments reported herein involved creation of aqueous NH2Cl solutions at a NH3-N: Cl molar ratio of 1.05:1. Interpretation of the apparent quantum yields for UVinduced photodecay of inorganic chloramines should take into account the effects of solution chemistry on observed behavior. For example, the generation of NCl3 will always occur in the presence of free chlorine (HOCl/OCl-). Although UV-induced photodecay of free chlorine compounds is a relatively efficient process, in terms of quantum yield, the rates of free chlorine photodecay are generally quite slow in UV systems, largely because HOCl and OCl- have low values of molar absorptivity for the wavelengths of interest in most UV photoreactors (12, 13). Nonetheless, photodecay of free chlorine has the potential to influence decay of NCl3. Therefore, the true quantum yields for NCl3 photodecay are likely to be smaller than the values reported in Table 1; however, insufficient information exists to determine the degree to which this is so. As described above, NHCl2 will undergo several “dark” reactions at near-neutral to slightly alkaline pH values that will increase the rate of its decay relative to decay that can be attributed to UV-induced photoreaction alone. Therefore, the true quantum yields for UV-induced photodecay of NHCl2 will also be somewhat smaller than the values reported in Table 1. Another important consideration in photodecay of inorganic chloramines is the effect of solution pH on the photochemical processes. To address this issue, chloramine solutions were prepared at pH values of 6.5, 7.5, and 8.5 and subjected to UV254 irradiation. These values of pH were chosen to be representative of the conditions that are likely to exist in water treatment applications where chloramination and UV irradiation are practiced together, including municipal drinking water production facilities and swimming pools. Figure 3 presents a graphical summary of the results of these experiments for all three inorganic chloramine compounds. The results indicate that the rate of photodecay of inorganic chloramines has little or no pH-dependence. In particular, the rate of monochloramine photodecay showed essentially no change across the pH range 6.5-8.5. With regard to di- and trichloramine, photodecay appeared to be slightly favored at elevated pH. Product Analysis. The products of inorganic chloramine photodecay were analyzed by both IC and MIMS. Table 2 summarizes the product yields for UV254 photodecay of inorganic chloramines for conditions corresponding to the experiments described by Figure 3. Under UV254 irradiation, the stable products resulting from photodecay of inorganic chloramines included nitrate, nitrite, nitrous oxide, and ammonium; of these, nitrate and nitrous oxide were common products for each of the chloramines. On the other hand, chloramine solutions are unstable in the dark at near-neutral pH, and are decomposed by a series of reactions that also result in the oxidation of ammoniacal-nitrogen and reduction of chlorine; however, the products of these “dark” reactions
FIGURE 3. UV254 photodecay of monochlormaine (top), dichloramine (middle), and trichloramine (bottom) as the function of UV dose at pH 6.5, 7.5, and 8.5. The secondary horizontal axis refers to the dark decay experiment with NHCl2. are different from those observed for UV254 photodecay (11). For example, when monochloramine decomposes in the dark, ammonium and nitrogen gas (N2) are the primary nitrogen decay products, with smaller quantities of nitrate (10). Under UV254 irradiation, nitrate, nitrite, and nitrous oxide were identified and quantified as the major products of photolysis of NH2Cl, accounting for the fate of more than 70% of the nitrogen initially present in the system (see Table 2). Watts and Linden reported nitrate and nitrite to be major products from photodecay of monochloramine; however, neutral gases (e.g., N2O or N2) and cations (e.g., NH4+) were not analyzed in their work (13). We also found that less than 20% of the original N present in solution was recovered as ammonium, and no N2 was formed during the UV254 photodecay of monochloramine. In general, the nitrogen-containing prodVOL. 43, NO. 1, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 2. Product Recovery from UV254 Photodecay of Inorganic Chloramines (Percentage of N Initially Present in the System As Chloramines)
SCHEME 1. Proposed reaction mechanism for UV photodecay of monochloramine
products chloramine precursor
pH
NO3-
NO2-
N2O
NH4+
total
NH2Cl
6.5 7.5 8.5
38.2% 9.7% 8.9%
10.8% 44.4% 56.2%
24.3% 23.6% 12.8%
16.8% 5.5% 14.7%
90.1% 82.2% 92.6%
NHCl2
6.5 7.5
15.4% 4.2%
30% 41%
45.4% 45.2%
NCl3
6.5 7.5 8.5
96% 65% 70%
2.3% 6% 9%
98.3% 71% 79%
ucts resulting from photodecay of monochloramine tended to be more highly oxidized than those produced by dark reactions. Nitrous oxide was identified as a common product of UV254 photodegradation of all three inorganic chloramine compounds, although the N2O yield was highly variable among the three inorganic chloramine compounds, and also showed pH dependence. Nitrous oxide has been reported to be a stable product resulting from UV irradiation of NH3 (14). Nitrous oxide has relatively low toxicity, but it is a major greenhouse gas. Although solution pH appears to have little or no effect on the rates at which inorganic chloramines photodecay, it does appear to influence the products of photodecay. The data presented in Table 2 indicate that nitrate is produced more efficiently in photodecay of monochloramine at low pH than other compounds, whereas a substantial fraction of N from monochloramine is converted to nitrite at high pH. N2O formation was suppressed at the highest pH condition used in these experiments, relative to the other pH conditions investigated. For di- and trichloramine, nitrate formation was also favored at low pH, whereas N2O formation from these compounds appeared to increase with pH. Nitrite and ammonium were not detected as UV254 photoproducts of dichloramine or trichloramine. Among the stable N-containing photoproducts listed above, nitrate and nitrite are present in natural waters, and both compounds have been associated with health problems. For example, the Maximum Contaminant Levels (MCLs) for nitrite and nitrate in drinking water are 1 mg/L (as N) and 10 mg/L (as N) as established by the U.S. EPA, and 0.5 mg/L and 50 mg/L as established by the EU, respectively. The results described above indicate that UV254 irradiation of water containing inorganic chloramines will increase the concentrations of these compounds in solution; the magnitude of this increase will depend on the initial chloramine concentration, solution pH, and the amount of UV to which the solution is exposed. Reaction Scheme. The UV photodecay of inorganic chloramines appears to be quite complex, particularly for dichloramine and trichloramine. Product recoveries from UV irradiation of monochloramine solutions were generally more complete than those of dichloramine and trichloramine. Therefore, the focus of efforts to describe inorganic chloramine photodecay mechanisms was on monochloramine; however, we expect the primary photochemical processes involving the three inorganic chloramines to be similar, whereas the subsequent reaction sequences may differ. The proposed mechanism for UV photodecay of monochloramine is illustrated in Scheme 1, and discussed below. 64
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The proposed mechanism was developed based on observed photoproducts and reaction kinetics, as well as previous reports of related reactions from the literature. The kinetics of chloramine photodecay depend on the initial reactant concentration. This suggests the involvement of UV-induced species in one or more rate-limiting steps. UV-induced cleavage of the N-Cl bond of monochloramine will yield an amine radical ( · NH2) and a chlorine radical (15) (step 1). The reaction of · NH2 with O2 in aqueous solution has been studied by several groups (16-18), and the results are consistent with a mechanism in which an intermediate peroxyl radical, NH2O2 · , is formed (step 2). Reported rate constants for this reaction range from 1.2 × 108 L mol-1 s-1 to 1.1 × 109 L mol-1 s-1. Once NH2O2 · is formed, it will decompose to stable products. Two parallel pathways of decomposition of this intermediate are proposed. Pathway 1 is based on the mechanism of Laszlo et al. (16), in which the peroxyl radical protonates, then decomposes to NO and H2O. Nitric oxide (NO) will be an intermediate with a short half-life in this process, because it can be easily oxidized to nitrite in oxygen-containing aqueous solution (19) (step 3). Although we did not observe the formation of 15 NO by MIMS, the lack of an NO signal may be attributed to rapid oxidation to nitrite. In turn, nitrite can be oxidized to nitrate by monochloramine (20, 21) (step 4). This reaction is general acid assisted, and nitrate formation is favored at low pH. This conclusion is consistent with the results listed in Table 2. At pH 6.5, nitrate and nitrite recoveries represented 38.2% and 10.8% of the initial N present in the system as NH2Cl. As pH increased, nitrite concentration in the resulting solution increased and nitrate decreased. At pH 8.5, only 8.9% of the initiate NH2Cl was recovered as nitrate. Pathway 2 involves the reaction of the aminylperoxyl radical with itself to yield HNO and H2O2 (22). HNO has been reported to dimerize very rapidly to yield nitrous oxide (23) (step 5). Nitrous oxide was also found to be one of the major products during the photodecay of monochloramine. The chlorine radical is recognized as a strong oxidant, and may react with other compounds by electron transfer and H-abstraction (24). In this case, chlorine radical may react with monochloramine to produce · NHCl radical by H-abstraction. NH2Cl + · Cl f · NHCl + HCl
(5)
Chorine radical can also react with water to produce · OH radical, and k(Cl + H2O) ) (2.5 ( 0.2) × 105 s-1 (25). The reaction between · OH radical and monochloramine is illustrated in eq 6, and the rate constant has been reported to be k ) (5.1 ( 0.6) × 108 L mol-1 s-1 (26). NH2Cl + · OH f · NHCl + H2O
(6)
Bicarbonate, carbonate, and phosphate are all known to be effective · OH radical scavengers: the rate constants were
reported to be k(OH + CO32-) ) 4.2 × 108 M-1 s-1; k(OH + HCO3-) ) 1.5 × 107 M-1 s-1; k(OH + HPO42-) ) 5 × 106 M-1 s-1 (27, 28). Photodecay of monochloramine was observed to slow down with the addition of bicarbonate buffer and phosphate buffer (see Figure S2). This result indicates that · OH radical is involved in this reaction, but it is not the major intermediate. · NHCl radical is similar to · NH2, in that it will combine with O2 to generate a peroxyl radical NHClO2 · , which has been measured by pulse radiolysis experiments with NH2Cl solutions saturated with O2 (26). The decay of NH2O2 · radical has been described above. The peroxyl radical, NHClO2 · , may follow a similar pathway to form NO and then nitrous acid or nitrite ions, then HNO and N2O.
Acknowledgments We are grateful to the National Swimming Pool Foundation (NSPF) for financial support of this research.
Supporting Information Available Detailed descriptions of the experimental setup, the development of eq 4 (used for quantum yield estimation), and information regarding the effects of carbonate and phosphate on chloramines photodecay. This material is available free of charge via the Internet at http://pubs.acs.org.
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